UNIT 1 · Principles of Chemistry

Ch 1 States of Matter

Chapter 1: States of Matter

Everything around you — solids, liquids and gases — is made of tiny particles you cannot see. This chapter explains the three states of matter in terms of how those particles are arranged, how they move, and how much energy they have, and how heating or cooling converts one state into another. It then looks at diffusion (particles spreading out), and at solutions, solutes and solvents, including how to measure and plot solubility.

States of Matter

Solids, liquids and gases are known as the three states of matter.
Think about these everyday facts, which the particle model explains:
You can't walk through a brick wall — the particles are held together by strong forces of attraction and can't move out of your way.
You can swim through water — you can push the liquid particles out of the way.
It is easy to move through a gas — there are (almost) no forces of attraction between the particles.
If you boil about 5 cm³ of water, the steam will fill an average bucket — gas particles spread out to fill all available space.
Most liquids are less dense than the solids they come from — gaps open up between particles when a solid melts.

The Arrangement of the Particles

📊 Key diagram — Figure 1.4: Particles in a solid, liquid and gas

SOLID                 LIQUID                 GAS
● ● ● ● ● ●          ● ● ● ● ●             ●   ●     ●
● ● ● ● ● ●          ●  ● ●  ●              ●     ●   ●
● ● ● ● ● ●          ● ● ● ● ●             ●   ●   ●
regular rows,         mostly touching,       far apart,
closely packed        random arrangement     random, moving fast

Solid — particles are usually arranged regularly and packed closely together.
In a solid, the particles can only vibrate about fixed positions; they cannot move around.
Strong forces of attraction between solid particles keep them together.
Liquid — particles are still mostly touching, but some gaps have appeared (this is why most liquids are less dense than their solids). The forces between particles are less effective, so particles can move around each other. The particles are arranged randomly.
Gas — particles are moving randomly at high speed in all directions; they are much further apart and there are (almost) no forces of attraction between them.
Kinetic energy (movement energy): particles in a solid have less kinetic energy than particles in a liquid, which have less kinetic energy than particles in a gas.

💡 Exam tip: When drawing a liquid, don't draw particles floating far apart — draw the particles mostly touching the particles next to them, but irregular/random. When drawing a gas, spread them out widely with no pattern.

Interconversions Between the Three States of Matter

Changing State Between Solid and Liquid

Melting: If you heat a solid, the energy from the heat source makes the particles vibrate faster and faster. Eventually they vibrate so fast that the forces of attraction can no longer hold them together; the particles can then move around each other — the solid melts to form a liquid.
Melting point — the temperature at which a solid melts. The particles in the liquid have more kinetic energy than those in the solid, so energy must be supplied to convert a solid into a liquid.
Freezing: If the liquid is cooled, the particles move around more and more slowly until the forces of attraction hold them in fixed positions and they pack closely together — the liquid freezes, forming a solid.
Freezing point — the temperature at which a liquid freezes.
⭐ The melting point and freezing point of a substance are exactly the same temperature — they are just different names depending on which direction the change is going.

Changing State Between Liquid and Gas

There are two different ways a liquid can become a gas: boiling and evaporation.

Boiling - Boiling occurs when a liquid is heated so strongly that the particles are moving fast enough to overcome all the forces of attraction between them. - Boiling only occurs at one temperature — the boiling point of the liquid. - The stronger the forces of attraction between particles, the higher the boiling point (more energy is needed to overcome those forces). - Condensing: If a gas is cooled, the particles eventually move slowly enough that forces of attraction between them start to form and hold them together as a liquid — the gas condenses.

Evaporation - In any liquid, the average speed of the particles varies, but at each temperature some particles move faster than the average and others more slowly. - Evaporation is when some very fast particles at the surface of the liquid have enough energy to overcome the forces of attraction, break away, and escape to form a gas. - Evaporation happens at any temperature — you see no bubbling; the liquid just slowly disappears if open to the air. - E.g. water evaporates from puddles even when the temperature is below 5 °C in a UK winter — the puddle water certainly does not boil at 5 °C, only at 100 °C. - Evaporation vs boiling: - Evaporation occurs at the surface of the liquid, at any temperature, with no bubbling. - Boiling occurs throughout the liquid, at one fixed temperature (the boiling point), with bubbling. - In a closed container, gas particles also collide with the surface of the liquid; if moving slowly enough they are held by attractive forces and become part of the liquid again. So in a closed container, evaporation and condensation both occur at the same time.

Changing State Between Solid and Gas: Sublimation

A small number of substances can change directly from a solid to a gas, or from a gas to a solid, at normal pressure, without involving any liquid in the process.
Sublimation — the conversion of a solid directly into a gas.
Deposition — the reverse process (gas directly to a solid); also called 'de-sublimation', and some people use the word 'sublimation' for both directions.
Example: carbon dioxide. At ordinary pressures there is no such thing as liquid carbon dioxide — it turns directly from a solid to a gas at −78.5 °C. Solid carbon dioxide is known as dry ice.

📊 Key diagram — energy transfers (Figures 1.5–1.8)

energy put in (HEAT)                       energy put in (HEAT)
   solid ⇄ liquid      (melting / freezing)    liquid ⇄ gas   (boiling / condensing)
energy taken out (COOL)                    energy taken out (COOL)

energy put in (HEAT)
   solid ⇄ gas  (sublimation / deposition)
energy taken out (COOL)

Working Out the Physical State of a Substance at a Particular Temperature

A substance is a solid at temperatures below its melting point; a liquid between its melting point and boiling point; and a gas above its boiling point.
You can decide the state at room temperature by comparing the melting and boiling points with room temperature.
Room temperature — in science usually taken to mean a temperature between 20 and 25 °C. There is not one fixed value, so when comparing changes of state near room temperature, always make clear which value of room temperature is being used.
Worked example from Figure 1.10 (a temperature line from −273 °C to 400 °C showing each substance's melting and boiling points):
Oxygen: room temperature is above its boiling point → oxygen is a gas at room temperature.
Bromine (melting point −7 °C, boiling point 59 °C): heating bromine from −100 °C:
- below −7 °C it is a solid at −100 °C;
- at −7 °C (its melting point) it becomes a liquid and stays liquid to 59 °C;
- room temperature lies between its melting and boiling points → bromine is a liquid at room temperature;
- above 59 °C bromine is a gas.
Lithium: its melting point is above room temperature → lithium is a solid at room temperature.

💡 Exam tip: Questions give melting/boiling point tables and ask for the state at a given temperature — quote the rule: below m.p. = solid; between m.p. and b.p. = liquid; above b.p. = gas.

Diffusion

Diffusion in Gases

Suppose someone releases smelly ammonia gas in the lab — within a minute or so, everyone can smell it. Particles in a gas are free to move around, but the spreading takes time.
At room temperature ammonia particles travel at about 600 metres per second, so they could cross the lab in less than 1/100 s (0.01 s) if they travelled in a straight line. But each particle keeps bouncing off air particles, so by the time the smell reaches every corner, each ammonia particle may have travelled 30 or more kilometres.
Diffusion — the spreading out of particles from where they are at a high concentration (lots of them in a certain volume) to where they are at a low concentration (fewer of them in a certain volume).

📊 Key diagram — Figure 1.12/1.13: Diffusion of bromine gas

   ┌─────────────────────┐
   │  air          lid   │   ← two gas jars, lids between them
   │  (colourless)       │
   ├─────────────────────┤
   │  bromine gas        │
   │  (brown vapour)     │
   └─────────────────────┘
Remove lids → brown colour spreads UP until both jars are uniformly brown
(air particles also diffuse downwards)

Demonstrating diffusion in gases: a lower gas jar contains brown bromine gas, the upper one air. When the lids are removed, the brown colour of the bromine diffuses upwards until both jars are uniformly brown (air also diffuses downwards) — particles move around at random to give an even mixture.
You can repeat the experiment with hydrogen and air, putting a lighted splint in at the end to find out where the gases have gone. People expect the much less dense hydrogen to all go to the top jar, but you actually get identical explosions from both jars — diffusion mixes them fully.

🔬 Practical safety — bromine diffusion demo (teacher demonstration): must be prepared in a working fume cupboard, wearing eye protection and chemical-resistant gloves. Inhalation of bromine by anyone with breathing difficulties may produce a reaction (possibly delayed) requiring urgent medical attention.

Showing That Particles of Different Gases Travel at Different Speeds

This experiment uses the reaction between ammonia (NH₃) and hydrogen chloride (HCl) gases to give white solid ammonium chloride (NH₄Cl):

NH₃(g) + HCl(g) → NH₄Cl(s)

Pieces of cotton wool are soaked in concentrated ammonia solution (a source of ammonia gas) and concentrated hydrochloric acid (a source of hydrogen chloride gas), and placed in the ends of a long glass tube with rubber bungs to stop the poisonous gases escaping.
Ammonia and hydrogen chloride particles diffuse along the tube; where they meet, a white ring of solid ammonium chloride forms.
The white ring takes time to form (the particles take time to diffuse along the tube).
The ring appears closer to the hydrochloric acid end of the tube.
Why: ammonia particles are lighter than hydrogen chloride particles and therefore move faster, so they travel further in the same time. (The relative molecular mass of ammonia is 17; that of hydrogen chloride is 36.5.)
Conclusion: lighter gas particles diffuse faster than heavier ones (at the same temperature).

🔬 Practical safety — ammonia/hydrogen chloride tube (teacher demonstration): requires eye protection, avoidance of skin contact and of inhaling fumes; the apparatus must be cleaned up in a working fume cupboard.

Diffusion in Liquids

Diffusion through a liquid is very slow if the liquid is completely still.
Example: a small jar of strongly coloured potassium manganate(VII) solution placed in a gas jar of water can take days for the colour to diffuse throughout all the water.
Reason: particles in a liquid move more slowly than particles in a gas, are much closer together, and there is less space for particles to move into without colliding with another one.

The Dilution of Coloured Solutions

Imagine dissolving 0.01 g of potassium manganate(VII) in 1 cm³ of water to give a deep purple solution. Taking 1 drop as 0.05 cm³, there are 20 drops per cm³, so each drop contains 0.0005 g.
If you dilute this solution with water until the total volume is 10 000 cm³, you should still just be able to see the purple colour — there are now 200 000 drops, so each drop must contain at least one 'particle' of potassium manganate(VII), meaning 0.01 g contains at least 200 000 'particles'.
In reality a single potassium manganate(VII) 'particle' weighs about 0.000 000 000 000 000 000 000 26 g — there are about 19 000 000 000 000 000 000 particles in 0.01 g! You need very large numbers of particles in each drop to see the colour.
Point of the calculation: particles are unimaginably small, and visible colour requires enormous numbers of them.
(Potassium manganate(VII) is an ionic compound containing more than one sort of particle — see Chapter 7.)

Solutes, Solvents and Solutions

When a solid dissolves in a liquid:
Solute — the substance that dissolves.
Solvent — the liquid the solute dissolves in.
Solution — the liquid formed when the solute dissolves in the solvent.
When you make a solution, the attractive forces between the particles in the solute (the solid) are being broken, while new attractive forces form between solvent particles and solute particles.
Whether a particular solid is soluble in any solvent depends on whether the new attractive forces are strong enough to overcome the old ones.
Only a certain amount of solute will dissolve in a fixed amount of solvent at a particular temperature.
Saturated solution — a solution which contains as much dissolved solid as possible at a particular temperature; there must be some undissolved solute present.

💡 Exam tip: For a solution to be genuinely saturated, undissolved solid must be present. If you heat a saturated solution, more solid can usually dissolve (for most solids) because solubility rises with temperature.

Measuring Solubility — CHEMISTRY ONLY

Solubility of a solid in a solvent at a particular temperature — the mass of solute which must dissolve in 100 g of solvent at that temperature to form a saturated solution. In other words, it is the maximum mass of solute that dissolves in 100 g of solvent at a particular temperature.
Units: g per 100 g of solvent.
Example: the solubility of sodium chloride (common salt) in water at 25 °C is about 36 g per 100 g of water.

🔬 Practical: Investigating the Solubility of a Solid in Water (potassium nitrate at 40 °C)

Apparatus: evaporating basin (weighed), boiling tube of water, thermometer, potassium nitrate, Bunsen burner, tripod, gauze.

Method: 1. Weigh an evaporating basin (e.g. 25.72 g). 2. Heat a boiling tube of water to just above 40 °C. 3. Add potassium nitrate to the water and stir rapidly until no more will dissolve and undissolved solid is left over. 4. Allow the solution to cool to exactly 40 °C. 5. Pour off some of the solution into the evaporating basin — it is important that you pour off only solution and no solid (you don't have to pour off all the solution). 6. Weigh the evaporating basin and contents. 7. Heat the evaporating basin and contents gently to evaporate off all the water. 8. When it looks as if all the water has evaporated, weigh the basin and contents again. 9. Heat again and re-weigh — repeat until the mass no longer changes. This is called heating to constant mass, and makes sure all the water really has evaporated.

Example results: | Measurement | Mass / g | |---|---| | Evaporating basin | 25.72 | | Evaporating basin + solution | 58.00 | | Evaporating basin + dry crystals | 38.00 |

Calculations:

mass of crystals = 38.00 − 25.72 = 12.28 g
mass of water    = 58.00 − 38.00 = 20.00 g

12.28 g dissolves in 20.00 g of water, so ×5 dissolves in 100 g:
12.28 × 5 = 61.4 g per 100 g of water

→ solubility of potassium nitrate at 40 °C = 61.4 g per 100 g of water

General equation:

solubility (g/100 g) = (mass of solute / mass of solvent) × 100

Safety: Wear eye protection; heat gently to avoid burns from hot solid 'spitting' out of the basin. If any solid spat out, a lower mass of solid would be recorded and the solubility would appear lower than the actual value.

💡 Exam tip: You must be able to turn raw masses into a solubility: mass of solute = (basin + crystals) − basin; mass of water = (basin + solution) − (basin + crystals). Then scale up to 100 g of water.

Extension (not required at International GCSE): some solutes can form supersaturated solutions — solutions containing more dissolved solid than you would expect at that temperature. Adding one tiny crystal makes all the extra solute crystallise out, leaving a normal saturated solution. (Undissolved solid present when you make a saturated solution prevents supersaturated solutions forming — don't worry about this at IGCSE.)

Solubility Curves — CHEMISTRY ONLY

The solubility of solids changes with temperature, and you can plot this on a solubility curve.
Most solids have solubility curves like those for potassium salts: solubility increases with temperature — either dramatically or just a little.

📊 Key diagram — Figure 1.18: Solubility curves for KNO₃ and KCl

solubility
(g per 100 g water)
  250 ┤                          KNO₃ (steep — rises dramatically)
  200 ┤                        ╱
  150 ┤                      ╱
  100 ┤                    ╱
   50 ┤                  ╱
      │                ╱   KCl (shallower — rises a little)
    0 ┼──────────────────────────────────
      0  10 20 30 40 50 60 70 80 90 100
                   temperature / °C

Using curves to find the mass of crystals formed when a saturated solution cools:
For potassium nitrate (KNO₃): at 90 °C, 200 g dissolves in 100 g water; at 30 °C, only 50 g dissolves.
So if a solution containing 200 g of KNO₃ in 100 g of water cools from 90 °C to 30 °C, 150 g of potassium nitrate must be released — it comes out as crystals. We say the potassium nitrate crystallises out (precipitates out) of the solution.

💡 Exam tip: Solubility varies with temperature, so always quote a temperature with a solubility value, e.g. "the solubility of sodium chloride at 30 °C is…". Lines of best fit on solubility graphs can be straight or curved — read the plotted points.

Worked example — potassium chloride (Chemistry only)

Solubility of potassium chloride at various temperatures:

Temperature / °C	10	30	40	70	90
Solubility / g per 100 g water	31.2	37.2	40.0	48.5	53.9

The curve is virtually a straight line.
(i) Solubility at 50 °C: read from graph ≈ 42.5 g per 100 g of water.
(ii) Maximum mass dissolving in 50 g of water at 40 °C: solubility at 40 °C is 40 g per 100 g, so half as much (20 g) dissolves in 50 g. General formula:

maximum mass that dissolves (g) = [mass of water (g) / 100 (g)] × solubility (g per 100 g)

E.g. with 34.6 g of water at 40 °C: (34.6 / 100) × 40 = 13.84 g of potassium chloride dissolves. - (iii) Temperature at which crystals first appear when cooling a solution containing 51.0 g of KCl in 100 g of water: crystals appear as soon as the solution becomes saturated. More than 51.0 g dissolves at 90 °C, but as the solution cools the solubility decreases; drawing a line across at 51.0 g shows this is the maximum mass that dissolves at 80 °C, so crystals first appear at temperatures below 80 °C. - (iv) Mass crystallising if cooled to 10 °C: solubility at 10 °C is 31.2 g per 100 g, so 31.2 g stays in solution. Started with 51.0 g, so:

mass of crystals = 51.0 − 31.2 = 19.8 g

19.8 g of potassium chloride crystallises out.

✅ Quick check

Name the six changes of state: solid → liquid, liquid → solid, liquid → gas (two possible names), gas → liquid, solid → gas, gas → solid.
A substance has melting point −7 °C and boiling point 59 °C. What state is it at −100 °C, at 25 °C, and at 100 °C?
In the ammonia/hydrogen chloride diffusion tube, why does the white ring form closer to the hydrochloric acid end?
6.00 g of solute dissolves in 25.0 g of water at 30 °C to form a saturated solution. What is the solubility in g per 100 g of water?
Explain, in terms of particles, why a solid has a fixed shape but a gas fills its container.

Answers

1. melting; freezing; boiling or evaporation; condensing; sublimation; deposition (also called 'de-sublimation'). 2. At −100 °C: solid (below melting point). At 25 °C: liquid (between −7 °C and 59 °C). At 100 °C: gas (above boiling point). 3. Ammonia particles are lighter (Mr = 17) than hydrogen chloride particles (Mr = 36.5), so they diffuse faster and travel further in the same time — the ring forms where they meet, nearer the HCl end. 4. (6.00 / 25.0) × 100 = **24 g per 100 g of water**. 5. In a solid the particles are closely packed, regularly arranged and held by strong forces of attraction, so they only vibrate about fixed positions — the solid keeps its shape. In a gas the particles are far apart with (almost) no forces of attraction, moving randomly at high speed in all directions, so they spread out to fill the container.

Ch 2 Elements, Compounds and Mixtures

Chapter 2: Elements, Compounds and Mixtures

Most substances in everyday life — air, food, drinks — are mixtures. This chapter distinguishes elements (one type of atom), compounds (elements chemically combined in fixed proportions) and mixtures (substances mixed without chemical reaction), and shows how their different properties allow separation. It then covers the key separation techniques: filtration, crystallisation, simple distillation, fractional distillation and paper chromatography — including how a chromatogram and Rf values are used to identify the components of a mixture.

Elements

Element — a substance that can't be split into anything simpler by chemical means; it contains only one type of atom.
In models or diagrams, elements are shown as atoms of a single colour or size.
Examples: a pure metal such as magnesium, oxygen gas, diamond (a form of carbon).
There are 118 elements, shown in the Periodic Table.
Most elements occur naturally, such as hydrogen, helium and sulfur; some have to be made artificially, such as einsteinium.
(Precise wording: it isn't completely true that elements consist of only one type of atom. A better way of saying it is that all the atoms in an element have the same atomic number — most elements are mixtures of isotopes, which have the same atomic number but different mass numbers due to different numbers of neutrons. Diagrams ignore these differences. Isotopes are discussed in Chapter 3.)

Compounds

Compound — a substance formed when two or more elements chemically combine.
The elements in a compound always combine in fixed proportions:
Hydrogen and fluorine always combine to form hydrogen fluoride, HF.
Magnesium and fluorine always combine to form magnesium fluoride, MgF₂ — the elements must combine in these ratios.
Other examples: carbon dioxide (CO₂) and methane (CH₄).
Diagrams of compounds show more than one type of atom bonded together.
Examples: water, silicon dioxide, sodium chloride.

Mixtures

Mixture — two or more substances mixed together with no chemical reaction occurring.
Mixtures can be made from elements and/or compounds.
Mixture of elements: nitrogen and oxygen (air).
Mixture of compounds: carbon dioxide and water (vapour).
Mixture of an element with a compound: carbon dioxide and nitrogen.
The various components can be in any proportion — for example, you can put any amount of sugar into your cup of tea or coffee (until it becomes saturated).

Simple Differences Between Mixtures and Compounds

Proportions

In water (a compound), every single water molecule has two hydrogen atoms combined with one oxygen atom — it never varies.
In a mixture of hydrogen and oxygen gases, the two could be mixed together in any proportion.
If you had iron metal and sulfur, you could mix them in any proportion you wanted; but in iron sulfide (FeS, a compound), the proportion of iron to sulfur is always exactly the same.

Properties

In a mixture of elements, each element keeps its own properties; the properties of the compound are quite different.
Example — iron and sulfur:
In a mixture: the iron is grey, the sulfur is yellow; the iron reacts with dilute acids (e.g. hydrochloric acid) to produce hydrogen, while the sulfur does not react with the acid.
The compound iron sulfide (FeS) reacts quite differently with acids — it produces poisonous hydrogen sulfide gas, which smells of bad eggs. (The reaction between iron sulfide and acids isn't needed for exam purposes at International GCSE.)
Example — hydrogen and oxygen: a mixture of hydrogen and oxygen is a colourless gas which explodes when you put a flame to it; the compound water is a colourless liquid which just puts out a flame.

Ease of Separation

Mixtures can be separated by physical means — methods that don't involve any chemical reactions, such as changing temperature or dissolving part of the mixture in a solvent such as water.
Example: a mixture of iron and sulfur is easy to separate using a magnet — the iron sticks to the magnet, the sulfur doesn't.
Example: cool a mixture of hydrogen and oxygen gases — oxygen condenses into a liquid at a much higher temperature than hydrogen (−183 °C as opposed to −253 °C), leaving liquid oxygen and hydrogen gas, which are easy to separate.
The elements in a compound cannot be separated by physical means. To convert iron sulfide into separate samples of iron and sulfur requires chemical reactions; to separate water into hydrogen and oxygen you must change it chemically using electrolysis (explained in Chapter 10).

Melting Point and Boiling Point

Pure substances (elements and pure compounds) melt and boil at fixed temperatures. For example, the melting point of water is 0 °C and its boiling point is 100 °C.
Mixtures usually melt or boil over a range of temperatures.
A mixture is not a pure substance.
Impurities lower the melting point of a substance and raise the boiling point:
Dissolving 10 g of common (table) salt (sodium chloride) in 1 litre of water lowers the melting point to about −0.6 °C and raises the boiling point to about 100.2 °C.
Impurity — a small amount of an unwanted substance in a sample that is otherwise pure.
The melting point is very useful for determining whether a substance is pure: record the temperature at which the sample starts to melt and the temperature at which it has fully melted. Pure aspirin is a white powder that melts at 138 °C. If the sample you made melts at 128–134 °C, it is quite impure because it melts over a wide range of temperatures (below the melting point of pure aspirin).

💡 Exam tip: Pure substance = fixed melting/boiling point. A substance melting over a range (e.g. 87–93 °C) is a mixture/impure; a substance melting sharply at a single temperature (e.g. 122 °C) is pure.

Separation of Mixtures

Separating mixtures is extremely important: in the processing of crude oil, in producing fresh water from salt water, in the enrichment of uranium, and in forensic science (components of a mixture usually have to be separated before they can be analysed).

Filtration

Filtration can be used to separate a solid from a liquid (e.g. sand from water), and to separate two solids if only one of them is soluble in water (see rock salt below).
Residue — the substance left in the filter paper (the solid).
Filtrate — the liquid that comes through the filter paper.

📊 Key diagram — Figure 2.6: Filtration

        ┌──────┐
        │funnel│
   ┌────┴───┐  │
   │ filter │  │   ← filter paper holds sand (residue)
   │ paper  │  │
   └────┬───┘  │
        │  ↓    │   ← water (filtrate) passes through
        └──────┘
        ┌──────┐
        │beaker│   ← filtrate collects here
        └──────┘

Crystallisation

Crystallisation can be used to separate a solute from a solution (e.g. sodium chloride from a sodium chloride solution).
Method: heat the solution in an evaporating basin to boil off some of the water until an almost saturated solution is formed. This can be tested by dipping a glass rod into the solution and seeing if crystals form quickly on its surface when removed. Turn off the Bunsen burner and allow crystals to form as more water evaporates and the solution cools. The crystals can then be removed from the mixture by filtration.

📊 Key diagram — Figure 2.7: Crystallisation

  sodium chloride solution        ← drop of solution sampled on glass rod
  in evaporating basin              (crystals form quickly on it when
      │                             near-saturated)
      │  gauze
      │  tripod
   ~~~~heat~~~~

🔬 Practical pathway — Making Pure Salt from Rock Salt

Rock salt consists of salt contaminated by various earthy or rocky impurities that are not soluble in water. 1. Crush the rock salt. 2. Mix with hot water — the salt dissolves, but the impurities don't. 3. Filter — the impurities are filtered off and remain on the filter paper. 4. The filtrate is now a salt solution. 5. Obtain the solid salt from the solution by crystallisation.
This is typical of separating any mixture of two solids, one of which is soluble in water and one of which isn't.

Simple Distillation

Simple distillation can be used to separate the components of a solution — for example, to collect pure water from a sodium chloride solution (crystallisation would only give the salt).
The solution is heated in a flask; the water boils and is condensed back to a liquid by the condenser; the salt remains in the flask. The sodium chloride solution eventually becomes so concentrated that the salt crystallises out.

📊 Key diagram — Figure 2.9: Simple distillation

        flask ─────────────┐
                            │
  sodium chloride solution  │  side-arm → condenser
        (heated)  ─────────┘   water OUT ┐
                            ┌────────────┤  ← cold water flows
                            │            │    around the condenser
                            └────────────┘  water IN ┘
                                          ↓
                                     pure water

💡 Exam tip — condenser water flow: water is always fed into the condenser at the lower end — that way it fills the condenser jacket better, and if the flow of water stops the jacket remains full of water.

Note the thermometer at the top of the flask records the boiling temperature of the vapour (liquids boil and condense at their boiling point, so the thermometer stays at the boiling point of the pure liquid while it distils).

Fractional Distillation

Fractional distillation is used to separate a mixture of liquids such as ethanol (alcohol) and water.
Ethanol and water are completely miscible with each other — you can mix them in any proportion and they form a single liquid layer.
You separate them by taking advantage of their different boiling points: water boils at 100 °C, ethanol at 78 °C.
Both liquids boil, but by careful heating you can control the temperature of the column so that all the water condenses in the column and trickles back into the flask; only the ethanol remains a vapour all the way to the top of the fractionating column and out into the condenser.

📊 Key diagram — Figure 2.10: Fractional distillation

       thermometer ──┐ ~78 °C
                     ↓
   fractionating     ┌───────────┐
   column (packed    │           │ condenser
   with glass        │           │   water OUT ┐
   beads)  ──────────┘           └──────────┬──┘
                                            ↓
   water/ethanol mixture              almost pure ethanol
   (heated in flask)                  (≈96% pure)

The fractionating column is often packed with glass beads (or something similar) to give a high surface area, which helps separation of the two vapours. (For reasons beyond International GCSE, an empty column works perfectly well in the lab, and it is impossible to remove the last 4% of water from the ethanol by distillation — the ethanol produced by this experiment is about 96% pure.)

💡 Exam tip: Simple distillation = one liquid from a solution (e.g. water from salt solution). Fractional distillation = separating two (or more) miscible liquids with different boiling points (e.g. ethanol 78 °C from water 100 °C).

Paper Chromatography

Paper chromatography can be used to separate a variety of mixtures; at International GCSE level it is usually used to separate mixtures of coloured inks or food colourings — most inks and food colourings are not just one colour but contain a mixture of dyes.
It can also separate mixtures of colourless substances (such as sugars), but then some method must be used to make the spots visible on the paper.
Chromatogram — the pattern of separated spots you get on the paper.
Two important ways the dyes differ, which causes separation:
their affinity for the paper (how well they 'stick' to the paper);
how soluble they are in the solvent which moves up the paper.
In Figure 2.11, spot C has hardly moved — either it is not very soluble in the solvent, or it has a very high affinity for the paper (or both). Spot A has moved almost as far as the solvent front — it must be very soluble in the solvent and not have much affinity for the paper.
If a dye does not move from the pencil line, it is not at all soluble in the solvent being used — find a different solvent. If a dye moves up with the solvent front, it is too soluble in that solvent — again, try a different solvent.

🔬 Practical: Investigating the Composition of Dye with Paper Chromatography

Apparatus: chromatography paper, pencil, teat pipette or capillary tube, beaker, small amount of solvent (water or a non-aqueous solvent), watch glass (lid).

Method: 1. Draw a line with a pencil across the chromatography paper, about 1 cm from the bottom. Do not use a pen — the colours in the ink may move up the paper with the solvent. 2. Put a spot (using a teat pipette or capillary tube) of the mixture of dyes on the pencil line and allow it to dry. 3. Suspend the chromatography paper in a beaker that contains a small amount of solvent so that the bottom of the paper goes into the solvent. It is important that the solvent is below the pencil line so that the inks/colourings don't just dissolve in the solvent. 4. Put a lid (such as a watch glass) on the beaker so the atmosphere becomes saturated with the solvent — this stops evaporation of the solvent from the surface of the paper. 5. When the solvent has moved up the paper to about 1 cm from the top, remove the paper and draw a pencil line to show where the solvent got to. - Solvent front — the highest level of the solvent on the paper at any time. 6. Leave the paper to dry so that all the solvent evaporates.

Choosing a solvent: water or a non-aqueous solvent (any solvent other than water) — which you use depends on what substances are present; a suitable solvent is usually found by experimenting with different ones.

Reading the chromatogram: The mixture in Figure 2.11 separated into three spots and must therefore have contained a minimum of three different dyes — it could contain more, because it is possible that one spot is made up of two coloured dyes that by coincidence moved the same distance. You could only confirm this by repeating the experiment with a different solvent.

🔬 Practical safety: Avoid skin contact with the solvents and dyes, especially if you have sensitive skin.

Using Paper Chromatography in Analysis

You can use paper chromatography to identify the particular dyes in a mixture.
Draw a pencil line on a larger sheet of paper and mark the original positions; place one spot of your unknown mixture and one spot of each known dye along the line. Develop the chromatogram as before.
The unknown mixture (m) has spots corresponding to dyes d1, d3 and d4 — they have the same colour as spots in the mixture AND have travelled the same distance. Although dye d2 is the same colour as one of the mixture's spots, it has travelled a different distance, so it must be a different compound.

Rf Values

Rf value — a ratio used to describe how far a spot moves. Rf stands for retardation factor.
Because the solvent (and therefore the spots) move different distances each time, we can't just report the distance moved by a spot — we work out a ratio instead:

Rf = distance moved by a spot (from the pencil line)
     ────────────────────────────────────────────────
     distance moved by the solvent front (from the pencil line)

Measure to the centre of the spot.
Example (Figure 2.13): for dye d3, Rf = 2.9 cm / 3.6 cm = 0.81.
Rf values of the mixture spots: blue spot Rf = 0.9/3.6 = 0.25; orange spot Rf = 2.0/3.6 = 0.56; green spot Rf = 2.9/3.6 = 0.81.
Rf values of the known dyes: d1 = 0.56, d2 = 0.36, d3 = 0.81, d4 = 0.25.
Because the spots in mixture m have the same Rf values as d1, d3 and d4, the mixture contains those dyes.
An Rf value must be between 0 and 1 — if you get a number bigger than 1 you have probably divided the numbers the wrong way round.
An Rf value has no units.
Rf values depend on the solvent used and on the type of paper. If the mixture and individual dyes are on separate pieces of paper, you can still compare Rf values as long as you use the same type of paper and the same solvent.

💡 Exam tip: A matching dye must match in both colour AND Rf value (or distance moved). Write the Rf formula as spot ÷ solvent front — a value above 1 means you swapped numerator and denominator.

✅ Quick check

Classify each as element, compound or mixture: hydrogen; sea water; magnesium oxide; blood; mud; calcium; honey.
A sample melts at 128–134 °C. Pure aspirin melts at 138 °C. What does this tell you, and why?
Which technique would you use to: (a) get potassium iodide from a potassium iodide solution; (b) get water from a potassium iodide solution; (c) separate ethanol from a mixture of ethanol and water; (d) separate a red dye from a mixture of red and blue dyes; (e) remove insoluble calcium carbonate from water?
A spot travels 3.2 cm and the solvent front travels 4.0 cm. Calculate the Rf value.
Why must the starting line in paper chromatography be drawn in pencil, not pen, and why must the solvent level be below that line?

Answers

1. Hydrogen — element; sea water — mixture; magnesium oxide — compound; blood — mixture; mud — mixture; calcium — element; honey — mixture. 2. It is impure — impurities lower the melting point and make it melt over a range of temperatures rather than at one fixed temperature. 3. (a) crystallisation; (b) simple distillation; (c) fractional distillation; (d) paper chromatography; (e) filtration. 4. Rf = 3.2/4.0 = **0.80** (no units). 5. Pen ink would dissolve in the solvent and travel up the paper, contaminating the chromatogram, so a pencil line is used instead. If the solvent level were above the pencil line, the spots would dissolve straight off the paper into the solvent instead of being carried up the paper by the rising solvent front.

Ch 3 Atomic Structure

Chapter 3: Atomic Structure

The 118 elements are the building blocks from which everything is made — from simple substances such as carbon to complex ones such as DNA. This chapter explores the nature of atoms and molecules, the sub-atomic particles (protons, neutrons and electrons) inside an atom, and how to describe an atom using its atomic number, mass number, isotopes and relative atomic mass (Ar), including calculating Ar from isotopic abundances.

Atoms and Molecules

Atom — the smallest piece of an element that can exist on its own and still be recognised as that element. Copper is an element: if you tried to cut it into smaller and smaller pieces, the final result would be the smallest possible piece of copper — an individual copper atom. You can split the atom into smaller pieces (protons, neutrons and electrons), but you would no longer have copper.
New atoms are produced in stars… or in nuclear processes such as nuclear bombs, nuclear reactors or radioactive decay.
Atoms can be joined together to make molecules.
Molecule — a group of two or more atoms chemically bonded (by covalent bonds). The atoms in a molecule can be from the same element or from different elements.
A hydrogen (H₂) molecule consists of 2 hydrogen atoms chemically bonded together.
A water (H₂O) molecule consists of 2 hydrogen atoms and an oxygen atom chemically bonded.
In diagrams of molecules, the lines between the atoms represent chemical bonds.

💡 Exam tip: "Element" = one type of atom; "molecule" = 2+ atoms held by covalent bonds (same or different elements); "compound" = molecule/ionic lattice made of 2+ different elements.

The Structure of the Atom

Atoms are made of protons, neutrons and electrons.
Sub-atomic particles — the particles smaller than an atom: protons, neutrons and electrons.
The nucleus of the atom contains the protons and neutrons.
The electrons are found at large distances from the nucleus (compared to the size of the nucleus); in a given case they are found most of the time somewhere in a certain region of space at some particular distance from the nucleus. Electrons with different energies are found at different distances from the nucleus.
Electrons are constantly moving — it is impossible to know exactly where an electron is at any moment; you can only identify that it has a particular energy and the region in which it is likely to be found. (Diagrams showing electrons orbiting the nucleus like planets around the sun can be misleading.)

📊 Key diagram — Figure 3.4: The structure of a helium atom

        (electron cloud — shaded pink region, most likely found here)
        . . . . . . . . . . . .
        .   ╭───────────────╮  .
        .  (   proton (+)   )  .     ● proton (charge +1)
        .  (   neutron 0    )  .     ○ neutron (charge 0)
        .  ╰───────────────╯  .     · electrons (charge −1) —
        . . . . . . . . . . . .        tiny, far out, constantly moving
          ← nucleus (tiny, dense) →

Scale: if you scaled a helium atom up to the size of a sports hall, the nucleus would be no more than the size of a grain of sand.

Table 3.1 — Properties of protons, neutrons and electrons:

Particle	Relative mass	Relative charge
proton	1	+1
neutron	1	0
electron	1/1836 (≈ 0.0005)	−1

Virtually all the mass of the atom is concentrated in the nucleus, because electrons have a much smaller mass than protons and neutrons.
The masses and charges are measured relative to each other because the actual values are incredibly small — it would take about 6 × 10²³ protons to weigh 1 g.

💡 Exam tip: You must learn Table 3.1 exactly: proton (+1, mass 1), neutron (0, mass 1), electron (−1, mass 1/1836). The electron is roughly 1840× lighter than a proton/neutron.

Atomic Number and Mass Number

Atomic number (also called the proton number) — the number of protons in an atom's nucleus.

atomic number = number of protons

The atomic number defines an element and is unique to that element: each of the 118 different elements has a different number of protons. If an atom has 8 protons it must be an oxygen atom.
You can identify an element by its atomic number instead of its name (e.g. 'a gold wristwatch' vs 'the element with atomic number 79'; 'the element with atomic number 17 is poisonous' instead of saying 'chlorine is poisonous').
Mass number (sometimes called the nucleon number) — the total number of protons and neutrons in the nucleus.

mass number = number of protons + number of neutrons

Notation — full symbol for an atom

        59            ← mass number (protons + neutrons), upper left
      Co
        27            ← atomic number (protons), lower left

This particular atom of cobalt contains 27 protons; to make the total up to 59, there must be 32 neutrons.
Therefore:

number of neutrons = mass number − atomic number

💡 Exam tip: The superscript (top) number is the mass number; the subscript (bottom) number is the atomic number. In two-letter symbols the first letter is a capital and the second is lower case — writing CO means carbon monoxide, not cobalt!

Isotopes

The number of neutrons in an atom of an element can vary slightly.
Example: there are three kinds of carbon atom — carbon-12, carbon-13 and carbon-14. They all have the same number of protons (all carbon atoms have 6 protons — that is their atomic number), but the number of neutrons varies (6, 7 or 8).
Isotopes — atoms of the same element which have the same atomic number (same number of protons) but different mass numbers (different numbers of neutrons).

📊 Key diagram — Figure 3.5: The nuclei of the three isotopes of carbon

   carbon-12           carbon-13           carbon-14
  6 protons           6 protons           6 protons
  6 neutrons          7 neutrons          8 neutrons
  ●●●●●●              ●●●●●●              ●●●●●●
  ○○○○○○              ○○○○○○○             ○○○○○○○○

The varying number of neutrons makes no difference to their chemical reactions: chemical properties are controlled by the number and arrangement of the electrons, and that is identical for all three isotopes.

Relative Atomic Mass

In a Periodic Table you might see, e.g., chlorine with the number 35.5. If you calculate the number of neutrons for chlorine: 35.5 − 17 = 18.5 — it is not possible to have half a neutron, so something must be wrong.
The number 35.5 is not the mass number for chlorine but rather the relative atomic mass (Ar): chlorine consists of two isotopes, ³⁵Cl and ³⁷Cl, and a naturally occurring sample contains a mixture of these.
Relative atomic mass (Ar) — the average mass of an atom, taking into account the amount of each isotope present (the isotopic abundances) in a naturally occurring sample of an element.
This type of average is called a weighted average (or weighted mean).
A naturally occurring sample of chlorine must contain more of the ³⁵Cl isotope than ³⁷Cl, because the relative atomic mass (35.5) is closer to 35 than to 37.
The number above each symbol in the International GCSE Periodic Table is a relative atomic mass, not a mass number. In most cases the relative atomic mass stated is the same as the mass number of the most common isotope; the only exceptions are chlorine (35.5) and copper (63.5).

Calculating Ar from isotopic abundances — worked example (boron)

A naturally occurring sample of boron contains 20% ¹⁰B and 80% ¹¹B. Calculate the relative atomic mass.

Imagine there are 100 atoms: 20 will have mass 10 and 80 will have mass 11.
Total mass of the 20 atoms with mass 10 = 20 × 10.
Total mass of the 80 atoms with mass 11 = 80 × 11.
Average = total mass ÷ 100 atoms:

         20 × 10 + 80 × 11
Ar(B) = ───────────────────  =  10.8
               100

The relative atomic mass of boron is 10.8.
Even if there are three or four different isotopes, do the calculation the same way: calculate the total mass of 100 atoms, then divide the answer by 100.

Ar = (mass₁ × abundance₁) + (mass₂ × abundance₂) + …  (abundances as %)
     ────────────────────────────────────────────────────────────────
                                  100

💡 Exam tip: Set the abundance percentages as numbers out of 100 (e.g. 78.99% → 78.99). Add the products and divide by 100. The answer should lie between the lightest and heaviest isotope masses — closer to whichever isotope is more abundant.

Counting the Number of Electrons in an Atom

Atoms are electrically neutral (no overall charge): the charge on a proton (+1) is equal but opposite to the charge on an electron (−1), so in an atom:

number of electrons = number of protons

If an oxygen atom (atomic number = 8) has 8 protons, it must also have 8 electrons; a chlorine atom (atomic number = 17) has 17 electrons.
The key feature is knowing the atomic number — you can find it from the Periodic Table (remember: the atomic number is the same as the number of protons).
The atomic number is defined in terms of protons, not electrons, because the number of electrons can change in chemical reactions — for example when atoms form ions (see Chapter 7).

The Periodic Table

Atoms are arranged in the Periodic Table in order of increasing atomic number.
Most Periodic Tables show two numbers against each symbol — be careful to choose the right one. The atomic number will always be the smaller number. The other number will be either the mass number of the most common isotope or the relative atomic mass of the element.
Using a Periodic Table to find the numbers of sub-atomic particles:

number of protons  = atomic number
number of electrons = number of protons
number of neutrons  = mass number − atomic number

Worked example — uranium

The symbol for uranium appears as:

238
U      (uranium, atomic number 92)
 92

The atomic number is the smaller number, so the atomic number of uranium is 92.
Atomic number = number of protons → uranium contains 92 protons.
Protons = electrons → uranium contains 92 electrons.
Neutrons = mass number − atomic number = 238 − 92 = 146 neutrons.

Common isotope questions (from chapter questions)

Fluorine atoms have a mass number of 19; from the Periodic Table its atomic number is 9 → 9 protons, 9 electrons, 19 − 9 = 10 neutrons. Protons must always equal electrons because atoms are electrically neutral.
Lithium has two isotopes, ⁶Li (abundance 7%) and ⁷Li (abundance 93%):

Ar(Li) = (6 × 7) + (7 × 93) / 100 = (42 + 651)/100 = 6.93

Magnesium has ²⁴Mg (78.99%), ²⁵Mg (10.00%) and ²⁶Mg (11.01%):

Ar(Mg) = (24×78.99 + 25×10.00 + 26×11.01)/100 = (1895.76 + 250 + 286.26)/100 = 24.32

Lead has ²⁰⁴Pb (1.4%), ²⁰⁶Pb (24.1%), ²⁰⁷Pb (22.1%), ²⁰⁸Pb (52.4%):

Ar(Pb) = (204×1.4 + 206×24.1 + 207×22.1 + 208×52.4)/100 = (285.6 + 4964.6 + 4574.7 + 10899.2)/100 = 207.24

Iridium has two isotopes ¹⁹¹Ir and ¹⁹³Ir; Ar = 192.22, which is closer to 193, so a naturally occurring sample contains more ¹⁹³Ir than ¹⁹¹Ir.
Considering only the most common isotope of each element, there is only one element with more protons than neutrons (hydrogen-1, 1 proton, 0 neutrons — you can check this with the Periodic Table), so the statement is true.

✅ Quick check

State the relative mass and charge of: proton, neutron, electron.
An atom of aluminium has a mass number of 27 and an atomic number of 13. How many protons, neutrons and electrons does it contain?
What are isotopes? Why do isotopes of the same element have identical chemical properties?
Chlorine has two isotopes: ³⁵Cl (75%) and ³⁷Cl (25%). Calculate the relative atomic mass of chlorine.
Why is the relative atomic mass of chlorine not a whole number, and would you expect a chlorine sample to contain more ³⁵Cl or ³⁷Cl?

Answers

1. Proton: relative mass 1, charge +1. Neutron: relative mass 1, charge 0. Electron: relative mass 1/1836, charge −1. 2. 13 protons, 13 electrons (equal to atomic number), neutrons = 27 − 13 = 14. 3. Isotopes are atoms of the same element with the same atomic number (same number of protons) but different mass numbers (different numbers of neutrons). Chemical properties are controlled by the number and arrangement of electrons, which is identical for all isotopes of an element. 4. Ar = (35 × 75 + 37 × 25)/100 = (2625 + 925)/100 = **35.5**. 5. Because natural chlorine is a mixture of two isotopes, ³⁵Cl and ³⁷Cl, and 35.5 is their weighted average — it is a weighted mean, not a mass number of a single isotope. There is more ³⁵Cl than ³⁷Cl because the average (35.5) is closer to 35 than to 37.

Ch 4 The Periodic Table

Chapter 4: The Periodic Table

The Periodic Table arranges all elements by atomic number, into groups (columns) and periods (rows). Elements in the same group have the same number of outer electrons and therefore similar chemical properties. This chapter covers group/period trends, metals vs non-metals, and the electron-shell patterns of the first 20 elements.

HOW THE PERIODIC TABLE IS ARRANGED

Element — a substance made of only ONE type of atom; cannot be broken down chemically.
Elements are arranged in order of increasing atomic number (number of protons).
Group — a vertical column (1–8/0). Elements in the same group have the same number of outer-shell electrons → similar chemical properties (e.g. Group 1 = 1 outer electron).
Period — a horizontal row. Elements in the same period have the same number of occupied electron shells.
KEY POINT: the group number = number of outer electrons (for Groups 1–7; Group 0/8 have a full outer shell).

ELECTRON SHELLS (first 20 elements)

Electrons occupy shells in order: 2, 8, 8, 2 (max 2 in the first shell, 8 in the next shells).
Sodium (Na, atomic number 11): 2.8.1 (1 outer electron → Group 1).
Chlorine (Cl, 17): 2.8.7 (7 outer electrons → Group 7).
Argon (Ar, 18): 2.8.8 (full outer shell → Group 0, unreactive).
Calcium (Ca, 20): 2.8.8.2 (2 outer electrons → Group 2).
💡 Exam tip: you must be able to draw the electron configuration (dot-and-cross or shell diagram) of the first 20 elements.

METALS AND NON-METALS

The staircase line divides the table: metals on the left, non-metals on the right.
Metals: shiny, conduct electricity and heat, malleable, high melting points, form positive ions (lose electrons), react with acids.
Non-metals: dull, poor conductors (except graphite), brittle, low melting points (often gases), form negative ions or share electrons (covalent).
Examples: metals — sodium, iron, copper, magnesium; non-metals — oxygen, chlorine, carbon, sulfur.
Metalloids (e.g. silicon) sit on the staircase — some properties of both.

GROUP TRENDS (key patterns)

Down a group: atoms get bigger (more shells), and the reactivity changes:
Group 1 (alkali metals): reactivity increases down the group.
Group 7 (halogens): reactivity decreases down the group.
Across a period: elements go from metallic to non-metallic.

THE GROUPS (see also Ch 11, 12)

Group 1 — Alkali metals (Li, Na, K): 1 outer electron; very reactive; soft; react with water to give hydrogen + hydroxide.
Group 2 — Alkaline earth metals (Mg, Ca): 2 outer electrons; less reactive than Group 1.
Group 7 — Halogens (F, Cl, Br, I): 7 outer electrons; reactive non-metals.
Group 0/8 — Noble gases (He, Ne, Ar): full outer shell → very unreactive; exist as single atoms (monatomic gases).

✅ Quick check

What does the group number tell you? (The number of outer electrons.)
What does the period number tell you? (The number of occupied electron shells.)
Draw the electron arrangement of sodium and chlorine. (Na 2.8.1; Cl 2.8.7.)
Why are noble gases unreactive? (Full outer shell of electrons.)
Which side of the table are metals on? (The left.)

Ch 5 Chemical Formulae, Equations and Calculations: Part 1

Chapter 5: Chemical Formulae, Equations and Calculations: Part 1

This chapter is about the language and maths of chemistry: writing formulae and equations, balancing them, and calculating relative atomic/formula mass, percentage composition, moles, and empirical formulae. (Part 2 covers gas volumes, concentrations and titrations.)

WRITING FORMULAE

Element symbols: each element has a symbol (H, O, C, Na, Cl, Fe...). The first letter is a capital, the second (if any) is lower case.
Formulae show the number of each atom (subscripts): H₂O = 2 H + 1 O; CO₂ = 1 C + 2 O; MgCl₂ = 1 Mg + 2 Cl.
Ions you must know: Na⁺, K⁺, Mg²⁺, Ca²⁺, Zn²⁺, Fe²⁺, Fe³⁺, Cu²⁺, Al³⁺, H⁺, NH₄⁺, Cl⁻, Br⁻, I⁻, OH⁻, NO₃⁻, SO₄²⁻, CO₃²⁻, O²⁻, S²⁻.
To write the formula of an ionic compound: the total positive charge must equal the total negative charge (swap and drop the numbers). E.g. Al³⁺ + SO₄²⁻ → Al₂(SO₄)₃; Ca²⁺ + OH⁻ → Ca(OH)₂.

WORD AND SYMBOL EQUATIONS

Word equation: magnesium + oxygen → magnesium oxide.
Symbol equation: 2Mg + O₂ → 2MgO.
Balancing: the number of each type of atom must be the same on both sides. Add coefficients (big numbers) — never change the subscripts.
State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).

RELATIVE MASSES

Relative atomic mass (Ar) — the average mass of an atom compared with 1/12th of a carbon-12 atom. (No units.) E.g. H = 1, C = 12, O = 16, Na = 23, Cl = 35.5.
Relative formula mass (Mr) — the sum of the Ar values of all the atoms in a formula.
Example: H₂O → Mr = (2×1) + 16 = 18.
CO₂ → 12 + (2×16) = 44.
MgCl₂ → 24 + (2×35.5) = 95.

THE MOLE

Mole (mol) — the amount of substance containing 6.02 × 10²³ particles (Avogadro's constant).
KEY equation: moles = mass (g) ÷ relative formula mass (Mr) n = m ÷ Mr
Example: how many moles in 36 g of water (Mr = 18)? n = 36 ÷ 18 = 2 mol.
Example: mass of 0.5 mol of CO₂ (Mr = 44) = 0.5 × 44 = 22 g.

PERCENTAGE COMPOSITION

Percentage by mass of an element = (total Ar of that element ÷ Mr of the compound) × 100%
Example: in CO₂ (Mr = 44), % C = (12 ÷ 44) × 100 = 27.3%; % O = (32 ÷ 44) × 100 = 72.7%.

EMPIRICAL FORMULA (simplest whole-number ratio)

Steps: 1. Write down the masses (or percentages) of each element. 2. Divide each by its Ar (relative atomic mass). 3. Divide all by the smallest result to get whole numbers. 4. These numbers are the formula's subscripts.
Example: a compound contains 2.4 g carbon and 6.4 g oxygen:
C: 2.4 ÷ 12 = 0.2; O: 6.4 ÷ 16 = 0.4.
Ratio 0.2 : 0.4 = 1 : 2 → CO₂.
To get the molecular formula, multiply the empirical formula by (Mr ÷ empirical formula mass).

REACTING MASS CALCULATIONS

Use the balanced equation: the mole ratio tells you the reacting amounts.
Example: 2Mg + O₂ → 2MgO. How much MgO from 4.8 g Mg?
Moles Mg = 4.8 ÷ 24 = 0.2 mol → moles MgO = 0.2 mol (1:1 ratio) → mass MgO = 0.2 × 40 = 8 g.

PERCENTAGE YIELD AND EXCESS REACTANTS

Percentage yield = (actual mass obtained ÷ theoretical mass expected) × 100% % yield = (actual ÷ theoretical) × 100
Yields are less than 100% because of: incomplete reactions (reversible), losses during filtration/transfer, side reactions.
Excess reagent: in a reaction, one reactant is often in excess (more than needed); the other is the limiting reactant — it is used up first and determines how much product forms. Use the moles of the limiting reactant for yield calculations.
Worked example: 10 g of calcium carbonate (Mr = 100) is heated: CaCO₃ → CaO + CO₂. Moles CaCO₃ = 10 ÷ 100 = 0.1 mol → theoretical CaO = 0.1 × 56 = 5.6 g. If only 4.2 g is collected, % yield = (4.2 ÷ 5.6) × 100 = 75%.

✅ Quick check

What is Ar and Mr? (Relative atomic mass; relative formula mass — sum of Ar.)
Write the formula of aluminium sulfate. (Al₂(SO₄)₃.)
Balance: Mg + O₂ → MgO. (2Mg + O₂ → 2MgO.)
How many moles in 80 g of NaOH (Ar: Na=23, O=16, H=1; Mr=40)? (80 ÷ 40 = 2 mol.)
What is the empirical formula of a compound with 4.8 g Mg and 3.2 g O? (Mg: 4.8/24=0.2; O: 3.2/16=0.2 → 1:1 → MgO.)

Ch 6 Chemical Formulae, Equations and Calculations: Part 2

Chapter 6: Chemical Formulae, Equations and Calculations: Part 2

⚠️ This entire chapter is flagged ✩ CHEMISTRY ONLY (it sits between a "CHEMISTRY ONLY" header and an "END OF CHEMISTRY ONLY" marker) — extended-tier content.

This chapter extends the mole idea from Chapter 5 into two new areas: gas volumes (using the molar volume of 24 dm³ at room temperature and pressure) and solutions (concentrations in g/dm³ and mol/dm³, plus acid–alkali titration calculations). Every calculation follows the same three-stage recipe: calculate moles of what you can, use the chemical equation to find moles of what you want, then convert to the required quantity.

Calculations involving gas volumes

Avogadro's law — equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
If you have 100 cm³ of hydrogen at some temperature and pressure, it contains exactly the same number of molecules as 100 cm³ of Cl₂ (or any other gas) under those same conditions, irrespective of the size of the molecules.
For helium and the other noble gases we would use the word 'atoms' instead of 'molecules' (noble gases exist as single atoms).
This works out nicely in reactions. Methane burns in oxygen:

CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

We need two O₂ molecules for each CH₄ molecule, so 100 cm³ of methane reacts exactly with 200 cm³ of oxygen, and 100 cm³ of carbon dioxide gas is formed — i.e. volumes react in the same ratio as the big numbers in the balanced equation.

📊 Key diagram (Fig 6.1/6.2): Three identical flasks containing different gases (e.g. Cl₂, NO, H₂) at the same temperature and pressure — all contain equal numbers of molecules. Fig 6.1: magnesium reacting with hydrochloric acid producing hydrogen gas (bubbles).

Units of volume

Volumes (of gases or liquids) are measured in:
cubic decimetres (dm³)
cubic centimetres (cm³)
litres (l) — the cubic decimetre is the preferred unit rather than the litre.
Conversion: 1 litre = 1 dm³ = 1000 cm³
To convert dm³ → cm³: multiply by 1000. To convert cm³ → dm³: divide by 1000.

1 dm³ = 1000 cm³      dm³ → cm³ : ×1000      cm³ → dm³ : ÷1000

The volume occupied by 1 mole of a gas

Molar volume — the volume occupied by 1 mole of a gas; the most important consequence of Avogadro's law.
1 mole of any gas occupies 24 dm³ (24 000 cm³) at rtp.
rtp — 'room temperature and pressure', usually taken as 20–25 °C and 1 atmosphere pressure.
The molar volume has true units of dm³/mol or cm³/mol, but at International GCSE you can treat it as a plain number (24 or 24 000) — just be careful with the units of volume.

📊 Key diagram (Figs 6.4 & 6.5): Mole–volume triangles — cover the quantity you want:

      volume of gas                volume of gas (cm³)
     ┌───────────┐                 ┌─────────────┐
     │  volume   │                 │   volume    │
     │ ┌───────┐ │                 │ ┌─────────┐ │
     │ │ moles │ │  =  mole ───── ├─│  moles  │─┤  24 000
     └─┴───────┴─┘     ratio      └─┴─────────┴─┘
        moles × 24                 moles × 24 000

volume of gas = number of moles × molar volume (use 24 for dm³ answers, 24 000 for cm³ answers)
number of moles = volume of gas ÷ molar volume

Calculations with molar volume

Calculating the volume of a certain number of moles

Worked example (book): Calculate the volume in dm³ of 0.20 mol CO₂ at rtp. Because we want the volume in dm³, use 24 dm³ as the molar volume: volume = number of moles × molar volume = 0.20 × 24 = 4.8 dm³

Calculating the volume of a given mass of gas

Worked example (book): Calculate the volume (in cm³) of 0.01 g of hydrogen at rtp (Aᵣ: H = 1). 1 mol H₂ has a mass of 2 g, so: number of moles = mass ÷ mass of 1 mol = 0.01/2 = 0.005 mol Because we want the volume in cm³, use the molar volume as 24 000 cm³: volume = 0.005 × 24 000 = 120 cm³

Calculating the number of moles from a volume

Worked example (book): Calculate the amount of moles in 120 cm³ of carbon dioxide. The volume is in cm³, so use 24 000 cm³ as the molar volume: number of moles = volume of gas ÷ molar volume = 120/24 000 = 0.005 mol

💡 Exam tip: Check the units of the gas volume before you start: if it is given in cm³ use 24 000; if it is given in dm³ use 24.

Calculations with chemical equations — using the molar volume

These follow the same method used for masses in Chapter 5, in three stages: 1. Calculate the number of moles of anything you can. 2. Use the chemical equation to deduce the number of moles of what you want. 3. Convert the number of moles to the required quantity (a mass or a volume).
Excess — more than enough of a reagent has been added to react with all of the other reactant (e.g. "an excess of dilute hydrochloric acid" means all the calcium carbonate reacts).

Worked example (book): Calculate the volume of carbon dioxide produced at rtp when an excess of dilute hydrochloric acid is added to 1.00 g of calcium carbonate. (Aᵣ: C = 12, O = 16, Ca = 40; molar volume = 24 dm³ at rtp.)

CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)

Moles of CaCO₃: Mᵣ = 40 + 12 + (3 × 16) = 100, so 1 mole has mass 100 g. moles = 1.00/100 = 0.0100 mol

From the equation, 1 mol CaCO₃ produces 1 mol CO₂, so 0.0100 mol CO₂.

volume of gas = moles × molar volume = 0.0100 × 24 = 0.24 dm³ of carbon dioxide.

💡 Exam tip: Include the units in your answer — the examiner will not know whether you mean cm³ or dm³ unless you write it down.

Worked example (book): Aluminium + acid — find the mass needed. What mass of aluminium would you need to add to an excess of dilute hydrochloric acid so that you produced 100 cm³ of hydrogen at rtp? (Aᵣ of Al = 27; molar volume = 24 000 cm³ at rtp.)

2Al(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2(g)

We can only calculate moles of the hydrogen (we don't know the mass of Al — that's what we want): moles of H₂ = 100/24 000 = 0.00417 mol (= 4.17 × 10⁻³ mol in standard form)

From the equation, 3 mol H₂ comes from 2 mol Al, so: moles of Al = ⅔ × 0.00417 = 0.00278 mol (Shortcut: divide the H₂ moles by 3 and multiply by 2, matching the 3 : 2 ratio.)

mass = number of moles × mass of 1 mol = 0.00278 × 27 = 0.075 g Al

💡 Exam tip: Multiply by 27 here, and not 54. Although the equation shows 2Al, you have already used the 2 when converting moles of H₂ to moles of Al. Also, the mass of 0.00278 mol of Al must always be the same — it cannot depend on what it is reacting with: 0.00278 mol of Al always contains the same number of atoms.

Concentrations of solutions

Concentration can be measured in either g/dm³ or mol/dm³ (also written g dm⁻³ and mol dm⁻³ — read as 'grams per cubic decimetre' and 'moles per cubic decimetre'). Remember: 1 cubic decimetre = 1 litre.
Molarity — the concentration of a solution in mol/dm³. A solution described as 2M means 2 mol/dm³, read as '2 molar'.
Converting between g/dm³ and mol/dm³ is no different from converting between moles and grams: it makes no difference that the substance is dissolved in 1 dm³ of solution.

Worked example (book): A sample of sea water has a concentration of sodium chloride of 35.1 g/dm³. Find its concentration in mol/dm³. (Aᵣ: Na = 23, Cl = 35.5)

1 mol NaCl has a mass of 58.5 g. number of moles = mass (g) ÷ mass of 1 mole (g) = 35.1/58.5 = 0.600 mol There are 0.600 mol in every dm³ of solution, so the concentration of NaCl is 0.600 mol/dm³.

Worked example (book) — Example 1: What is the concentration of a 0.050 mol/dm³ solution of sodium carbonate, Na₂CO₃, in g/dm³? (Aᵣ: C = 12, O = 16, Na = 23)

1 dm³ of solution contains 0.050 mol Na₂CO₃. 1 mol Na₂CO₃ weighs 106 g (Mᵣ = 2×23 + 12 + 3×16 = 106). 0.050 mol weighs 0.050 × 106 = 5.3 g, so the concentration is 5.3 g/dm³.

Worked example (book) — Example 2: What is the concentration in mol/dm³ of a solution containing 2.1 g of sodium hydrogencarbonate, NaHCO₃, in 250 cm³ of solution? (Aᵣ: H = 1, C = 12, O = 16, Na = 23)

1 mol NaHCO₃ has a mass of 84 g. 2.1 g is 2.1/84 = 0.025 mol. This is in 250 cm³, but we need the amount in 1 dm³ (1000 cm³) — there are 4 lots of 250 cm³ in 1000 cm³. Each portion contains 0.025 mol, so there must be 4 × 0.025 = 0.10 mol in 1000 cm³. The concentration is 0.10 mol/dm³.

📊 Key diagram (Fig 6.8): Solution triangle

        ┌─────────┐
        │  moles  │
        │ ┌─────┐ │
        │ │conc │ │  moles = volume (dm³) × concentration (mol/dm³)
        └─┴─────┴─┘
          vol × C

Working out a number of moles from a volume and a concentration

number of moles = volume of solution (dm³) × concentration (mol/dm³)
Because concentration is usually in mol/dm³ but volumes of solutions are often given in cm³, you usually have to convert the volume from cm³ to dm³ by dividing by 1000.

Worked example (book) — Example 3: Calculate the number of moles of NaOH in 50 cm³ of 0.10 mol/dm³ solution. Converting the volume to dm³: 50/1000 = 0.050 dm³. Number of moles = volume (dm³) × concentration = 0.050 × 0.10 = 0.0050 mol

A calculation involving hard water — calculations with equations involving solutions

Limescale (calcium carbonate, CaCO₃) can be removed from kettles by reacting it with a dilute acid such as ethanoic acid (present in vinegar):

CaCO3(s) + 2CH3COOH(aq) → (CH3COO)2Ca(aq) + CO2(g) + H2O(l)

Worked example (book): What mass of calcium carbonate can be removed by 50 cm³ of a solution of ethanoic acid that has a concentration of 2 mol/dm³? (Aᵣ: C = 12, O = 16, Ca = 40)

Moles of ethanoic acid = volume (dm³) × concentration = (50/1000) × 2 = 0.1 mol

From the equation there is a 2 in front of the CH₃COOH but no number (meaning 1) in front of the CaCO₃: 2 mol of CH₃COOH react with 1 mol of CaCO₃, so 0.1 mol of ethanoic acid reacts with 0.1/2 = 0.05 mol of CaCO₃.

mass = moles × mass of 1 mol = 0.05 × 100 = 5 g of calcium carbonate

💡 Exam tip: Don't be scared by unfamiliar chemicals (like CH₃COOH) — as long as you realise it is ethanoic acid, that's all you need to worry about for the calculation.

Acid-alkali titrations — calculations from titrations

Titration — a technique used to find out how much of one solution reacts with a certain volume of another solution of known concentration.
Method: a solution of an alkali is measured into a conical flask using a pipette. An acid is run in from a burette, swirling the flask constantly. Towards the end, the acid is run in a drop at a time until the indicator just changes colour (the endpoint; Fig 6.9 shows the endpoint using methyl orange).
If you know the concentration of either the acid or the alkali, you can use the titration results to find the concentration of the solution you reacted it with.
The acid and alkali can also be added the other way round (alkali in the burette).
Acid–alkali titrations are discussed in more detail in Chapters 16 and 17.

The standard calculation

Worked example (book): 25.00 cm³ of 0.100 mol/dm³ sodium hydroxide solution required 23.50 cm³ of dilute hydrochloric acid for neutralisation. Calculate the concentration of the hydrochloric acid.

NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

Moles of NaOH = (25.00/1000) × 0.100 = 0.00250 mol

From the equation, 1 mol NaOH reacts with 1 mol HCl, so 0.00250 mol NaOH reacts with 0.00250 mol HCl.

Convert moles to concentration: the 0.00250 mol of HCl was in 23.50 cm³ = 0.02350 dm³. concentration (mol/dm³) = number of moles (mol) ÷ volume (dm³) = 0.00250/0.02350 = 0.106 mol/dm³

💡 Exam tip: Remember to divide the volume by 1000 to convert to dm³ before using it.

A very slightly harder calculation (2:1 ratio)

Worked example (book): 25.0 cm³ of sodium hydroxide solution of unknown concentration was titrated with dilute sulfuric acid of concentration 0.050 mol/dm³. 20.0 cm³ of the acid was required to neutralise the alkali. Find the concentration of the sodium hydroxide in mol/dm³.

2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)

Moles of sulfuric acid used = (20.0/1000) × 0.050 = 0.0010 mol

The equation shows 1 mol of sulfuric acid reacts with 2 mol of sodium hydroxide, so the number of moles of sodium hydroxide = 2 × 0.0010 = 0.0020 mol. (That's what makes this different from the 1:1 example.)

That 0.0020 mol was in the 25.0 cm³ (25/1000 = 0.025 dm³) of NaOH solution: concentration = moles ÷ volume = 0.0020/0.025 = 0.080 mol/dm³

Reversing the calculations — finding a volume

Worked example (book) — Example 4: Calculate the volume of 0.100 mol/dm³ sodium hydrogencarbonate (NaHCO₃) solution needed to neutralise 20.0 cm³ of 0.125 mol/dm³ hydrochloric acid (HCl).

NaHCO3(aq) + HCl(aq) → NaCl(aq) + CO2(g) + H2O(l)

We have the volume and concentration of the HCl (and not enough information about anything else), so we must start here: moles of HCl = (20.0/1000) × 0.125 = 0.00250 mol

The equation shows the same number of moles of sodium hydrogencarbonate is needed: 0.00250 mol of NaHCO₃.

Rearrange the concentration equation: volume (dm³) = number of moles (mol) ÷ concentration (mol/dm³) volume = 0.00250/0.100 = 0.0250 dm³ (dm³ because the concentration is in mol/dm³) Convert to cm³ by multiplying by 1000: 25.0 cm³ of the sodium hydrogencarbonate solution.

📊 Key diagrams (Figs 6.9): titration flask at the endpoint with methyl orange — a single drop of acid has just turned the indicator's colour.

✅ Quick check

Calculate the volume in dm³ of 0.10 mol of gas at rtp (molar volume 24 dm³).
How many moles are there in 4800 cm³ of a gas at rtp?
Calculate the mass of CaCO₃ (Mᵣ = 100) that reacts with 25.0 cm³ of 2.00 mol/dm³ hydrochloric acid, given CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g).
20.0 cm³ of 0.100 mol/dm³ sodium hydroxide was neutralised by 25.0 cm³ of nitric acid. Calculate the concentration of the nitric acid (NaOH + HNO₃ → NaNO₃ + H₂O).
State the volume occupied by 1 mole of any gas at rtp.

Answers

1. volume = 0.10 × 24 = **2.4 dm³** 2. 4800/24 000 = **0.20 mol** 3. moles HCl = (25.0/1000) × 2.00 = 0.050 mol; moles CaCO₃ = 0.050/2 = 0.025 mol; mass = 0.025 × 100 = **2.5 g** 4. moles NaOH = (20.0/1000) × 0.100 = 0.0020 mol; 1:1 so moles HNO₃ = 0.0020 mol; concentration = 0.0020/(25.0/1000) = 0.0020/0.025 = **0.080 mol/dm³** 5. **24 dm³ (24 000 cm³)** at room temperature and pressure (rtp)

Ch 7 Ionic Bonding

Chapter 7: Ionic Bonding

A compound forms when two or more elements chemically combine. This chapter covers one way elements combine: transfer of electrons to form ionic compounds. Sodium (a dangerously reactive metal stored under oil) and chlorine (a poisonous, reactive gas) combine to make sodium chloride — salt, which is safe to eat in small quantities: the properties of a compound are very different from those of its elements. You will learn how ions form, how to draw dot-and-cross diagrams, how to work out ionic formulae, and why ionic substances have the physical properties they do.

Ionic bonding

Ionic compound — a compound held together by ionic bonding. Sodium chloride is the best-known example; others include magnesium oxide, calcium fluoride and zinc bromide.
All these compounds contain a metal combined with a non-metal — you can recognise ionic compounds because they (usually) contain a metal.
One or two exceptions exist: ionic compounds containing the ammonium ion (NH₄⁺) contain no metal, e.g. NH₄Cl and (NH₄)₂SO₄ (looked at later).
When a non-metal such as chlorine combines with a metal such as sodium, the chlorine atom has a stronger attraction for electrons than the sodium atom, and an electron is transferred from the outer shell of the sodium atom to the outer shell of the chlorine atom.
Because an electron has a negative charge: when something loses an electron it becomes positively charged; when something gains an electron it becomes negatively charged.
Sodium (atomic number 11) has 11 protons (11+ on the nucleus) and 11 electrons in the atom (11−), so no overall charge. After losing one electron the sodium ion has 10 electrons, so with 11+ and 10− there is an overall charge of 1+.
Ions — charged particles formed when atoms (or groups of atoms) lose or gain electrons. Ions can have either a positive or a negative charge.
Cation — a positive ion, for example Na⁺.
Anion — a negative ion, for example Cl⁻.
Ionic bonding — the strong electrostatic attraction between positive and negative ions. (This is a simplification — in reality you don't react sodium with chlorine atoms but with chlorine molecules, Cl₂.)
Electrostatic attraction — simply means that positively charged particles attract negatively charged particles.
When an ionic compound is formed, electron(s) are transferred from a metal atom to a non-metal atom to form positive and negative ions; the compound is held together by these attractions.
Ionic bonds are usually formed only if small numbers of electrons need to be transferred — typically 1 or 2, but occasionally 3.

Dot-and-cross diagrams

Ionic bonding is often shown using dot-and-cross diagrams.
Although electrons are drawn as dots or as crosses, there is absolutely no difference between them in reality — the dots and crosses simply show that the electrons have come from two different atoms. You could equally well use two different coloured dots or two different coloured crosses.

📊 Key diagram (Fig 7.2): Ionic bonding in sodium chloride — a sodium atom (with its single outer electron drawn as a dot) transfers that electron to a chlorine atom (outer shell drawn with crosses), forming Na⁺ (2,8) and Cl⁻ (2,8,8). Fig 7.3: part of a sodium chloride crystal held together by attraction between positive and negative ions.

Ionic bonding in magnesium oxide

📊 Key diagram (Fig 7.4): Dot-and-cross for magnesium oxide — two electrons (dots) are transferred from a magnesium atom to an oxygen atom, forming Mg²⁺ and O²⁻ ions.

2 electrons are transferred from a magnesium atom to an oxygen atom to form Mg²⁺ and O²⁻ ions.

💡 Exam tip: When drawing dot-and-cross diagrams, don't forget the charges on the ions.

The significance of noble gas electronic configurations in ionic bonding

Looking at the ions in Figures 7.2, 7.4, 7.5 and 7.6, each has a noble gas electronic configuration: (2, 8) — the same as neon — or [2, 8, 8] — the same as argon (for the first 20 elements).
Atoms lose or gain electrons so that they achieve a noble gas electronic configuration.
Elements in Groups 1, 2 and 3 lose their outer shell electrons to form 1+, 2+ and 3+ ions.
Elements in Groups 5, 6 and 7 gain electrons to form 3−, 2− and 1− ions.
Isoelectronic — means 'have the same number of electrons'. When sodium forms an ion (electronic configuration 2, 8) it becomes isoelectronic with neon [2, 8]. Na⁺ and Ne are isoelectronic.
Elements in Groups 1, 2, 6 and 7 always form ions that are isoelectronic with the nearest noble gas atom:
Rubidium (Rb, Group 1) loses its 1 outer-shell electron to form a 1+ ion with the same number of electrons as a krypton (Kr) atom.
Iodine (Group 7) gains 1 electron to form a 1− ion with the same number of electrons as a xenon (Xe) atom.
However, many common ions do not have noble gas structures: Fe²⁺, Fe³⁺, Cu²⁺, Zn²⁺, Ag⁺ and Pb²⁺ are all encountered on the International GCSE course — not one of them has a noble gas structure (you won't have to write their electronic structures).

📊 Key diagrams: Fig 7.5 — an aluminium atom loses its 3 outer-shell electrons to form an Al³⁺ ion; Fig 7.6 — a nitrogen atom gains 3 electrons to form the nitride ion (N³⁻); Fig 7.7 — Na⁺ and Ne are isoelectronic; Fig 7.8 — an Fe²⁺ ion is definitely not a noble gas structure.

Other examples of ionic bonding

Lithium fluoride

📊 Key diagram (Fig 7.9): Dot-and-cross for lithium fluoride — 1 electron is transferred from the lithium atom to the fluorine atom, forming Li⁺ and F⁻.

The lithium atom has 1 electron in its outer shell that is easily lost; the fluorine atom has space to receive one. One electron is transferred from lithium to fluorine. Lithium fluoride is held together by strong electrostatic attractions between positive Li⁺ ions and negative F⁻ ions.

Calcium chloride

📊 Key diagram (Fig 7.10): Dot-and-cross showing the formation of calcium chloride — the calcium atom [2, 8, 8, 2] transfers its 2 outer electrons, one to each of two chlorine atoms [2, 8, 7].

The calcium atom [2, 8, 8, 2] has 2 electrons in its outer shell, but each chlorine atom [2, 8, 7] only has room in its outer shell to take one of them — so you need two chlorines for every calcium. The 2 electrons are transferred from the calcium atom to two chlorine atoms, one to each.
The formula of calcium chloride is therefore CaCl₂.
There are very strong electrostatic attractions holding the ions together because of the 2+ charge on the calcium ion.

Formulae for ionic compounds

There are so many ionic compounds at International GCSE that it would be impossible to learn all their formulae — you need a simple shortcut method. (You could work a few out from first principles using electronic structures, but that would take ages; others would be too difficult.)

The need for equal numbers of pluses and minuses

Ions carry an electrical charge, either positive or negative.
Compounds are electrically neutral, so in an ionic compound there must be the right number of each sort of ion so that the total positive charge exactly cancels out the total negative charge.
To work out a formula you therefore need to know the charges on the ions.

Cases where you can work out the charge on an ion

Any element in Group 2 has 2 outer electrons, which it will lose to form a 2+ ion.
Any element in Group 6 has 6 outer electrons and has room to gain 2 more; this leads to a 2− ion.
Similar arguments apply in the other groups:

Group in Periodic Table	Charge on ion	Example
1	1+	Na⁺
2	2+	Mg²⁺
3	3+	Al³⁺
5	3−	N³⁻
6	2−	O²⁻
7	1−	Br⁻

All metals form positive ions. Elements in Group 4 only form a few ionic compounds and the situation is more complicated; you will need to learn that lead forms a 2+ ion (Pb²⁺).
You will always have a copy of the Periodic Table, even in an exam, so you can always find out which group an element is in.

Cases where the name tells you the charge

All metals form positive ions. Names such as lead(II) oxide, iron(III) chloride or copper(I) sulfate tell you directly about the charge on the metal ion — the number after the metal tells you the number of charges:
lead(II) oxide contains a Pb²⁺ ion
iron(III) chloride contains an Fe³⁺ ion
copper(I) sulfate contains a Cu⁺ ion
Some ion charges cannot be worked out — you have to learn them:

Positive ion	Charge	Substance	Negative ion	Charge	Substance
Zn²⁺	2+	zinc	NO₃⁻	1−	nitrate
Ag⁺	1+	silver	OH⁻	1−	hydroxide
H⁺	1+	hydrogen	CO₃²⁻	2−	carbonate
NH₄⁺	1+	ammonium	SO₄²⁻	2−	sulfate

These are the difficult ones — be sure to learn both the formula and the charge for each ion.
Ammonium chloride (NH₄Cl) is an example of an ionic compound that does not contain a metal: there is ionic bonding between the NH₄⁺ and Cl⁻ ions, but the NH₄⁺ ion itself is held together by covalent bonding (see Chapter 8).

Confusing endings!

Don't confuse ions such as sulfate and sulfide.
Any 'ide' ending means there isn't anything complicated — just the two elements named, e.g. sodium chloride is sodium and chlorine combined. Copper(I) sulfide (Cu₂S) contains Cu⁺ and S²⁻ ions only.
Once you have an 'ate' ending, there is oxygen (and possibly other things) there as well: copper(II) sulfate contains copper, sulfur and oxygen.

💡 Exam tip: Not looking carefully at word endings is one of the most common mistakes students make when they start to write formulae. Be careful!

📊 Key diagrams: Fig 7.11 — copper(I) sulfide is Cu₂S; Fig 7.12 — copper(II) sulfate is CuSO₄.

Deducing the formula for an ionic compound

Example 1: Sodium oxide

Sodium is in Group 1, so the ion is Na⁺.
Oxygen is in Group 6, so the ion is O²⁻.
To balance the two negative charges on one oxide ion you need two sodium ions:

Na⁺  Na⁺  O²⁻   →   formula: Na2O

Example 2: Barium nitrate

Barium is in Group 2, so the ion is Ba²⁺.
Nitrate ions are NO₃⁻ (you have to remember this).
You need two nitrate ions for each barium ion to balance the charges.

Ba²⁺  2 × NO3⁻   →   formula: Ba(NO3)2

💡 Exam tip (brackets): Brackets must be written if you have more than one of these complex ions (ions containing more than one atom). Without the brackets the formula would look like BaNO₃₂ — which would read as 1 barium, 1 nitrogen and 32 oxygens! In any other situation brackets are completely unnecessary.

Example 3: Iron(III) sulfate

Iron(III) tells you the metal ion is Fe³⁺.
Sulfate ions are SO₄²⁻.
To have equal numbers of positive and negative charges you need two iron(III) ions for every three sulfate ions, giving 6+ and 6− in total.

Fe³⁺  SO4²⁻   →   formula: Fe2(SO4)3

The shortcut: crossing over the numbers in the charges

📊 Key diagrams (Figs 7.13 & 7.14): Cross-over method — a shortcut for complicated formulae: swap over the numbers in the charges.

       3+     2−
   Fe³⁺   SO4²⁻   →   Fe2(SO4)3
   (the 3 drops onto SO4, the 2 drops onto Fe)

       2+     1−
   Ca²⁺    Cl⁻   →   CaCl2
   (no extra number in front of Cl because we don't tend to write in a 1)

For calcium chloride, cross over the numbers in the charges — but there is no number written in front of Cl because we do not tend to write in a 1.
Ion charges are not usually shown in formulae. They can be shown (e.g. Na⁺Cl⁻ is sometimes written to make a particular point), but for most purposes the charges are left out — in an ionic compound the charges are there whether you write them or not.

⚠️ Caution: You can get the wrong answer when the charges on the ions are the same. For example, the formula of calcium oxide is CaO and not Ca₂O₂. When the charges on the positive and negative ions are the same, there will be 1 of each ion in the formula, so there is no need to swap anything over.

📊 Key diagram (Fig 7.15): calcium oxide is CaO, not Ca₂O₂.

Giant ionic structures

Lattice — a regular array of particles (like a lattice fence — a regular, repeating structure).
All ionic compounds form crystals that consist of lattices of positive and negative ions packed together in a regular way.
The lattice is held together by the strong electrostatic attractions between the positively and negatively charged ions.

The structure of sodium chloride

📊 Key diagram (Figs 7.17 & 7.18): Structure of sodium chloride — a model of a small part of a sodium chloride crystal, and an 'exploded' view showing that each Na⁺ ion is surrounded by 6 Cl⁻ ions and each Cl⁻ ion is surrounded by 6 Na⁺ ions. The lines in the exploded diagram are not bonds — they are just there to show the arrangement; ions joined by lines are touching each other.

Each sodium ion is surrounded by 6 chloride ions; in turn each chloride ion is surrounded by 6 sodium ions. This pattern repeats itself throughout the structure over vast numbers of ions.
The structure of sodium chloride is described as a giant ionic lattice.
'Giant' here is not used in the sense of 'big' — it describes a structure in which there are no individual molecules. All the sodium ions in the lattice attract all the chloride ions; we cannot pick out sodium chloride molecules. The lattice extends throughout the structure in all directions, with no limit to the number of sodium and chloride ions — all we know is that there must be the same number of each.

💡 Exam tip (very important): You must not talk about molecules of an ionic compound. There are no individual molecules. Writing "molecules" in the exam will be marked wrong and you could lose all the marks for a question.

The structure of magnesium oxide

Magnesium oxide, MgO, contains Mg²⁺ and O²⁻ ions. It has exactly the same structure as sodium chloride.
The only difference: the magnesium oxide lattice is held together by stronger forces of attraction, because in MgO, 2+ ions are attracting 2− ions. In sodium chloride the attractions are only between 1+ and 1− ions.
The Mg²⁺ ion is also smaller than the Na⁺ ion, and O²⁻ is smaller than Cl⁻ — this causes stronger attractions, but the effect of the charge on the ions is more important.

The physical properties of ionic substances

High melting points and boiling points — because of the strong electrostatic forces of attraction holding the lattice together. A lot of energy has to be supplied to break the strong electrostatic forces of attraction between oppositely charged ions in the giant lattice structure.
Crystalline — ionic compounds tend to be crystalline, reflecting the regular arrangement of ions in the lattice. Sometimes the crystals are too small to be seen except under powerful microscopes — magnesium oxide, for example, is always seen as a white powder because its individual crystals are too small to see with the naked eye.
Brittle — ionic crystals tend to be brittle. Any small distortion of a crystal brings ions with the same charge alongside each other; like charges repel and so the crystal splits itself apart.

📊 Key diagram (Fig 7.20): Why ionic crystals are brittle — an applied force distorts the crystal so that ions of the same charge are pushed alongside each other; the repulsion between ions with the same charge forces the crystal apart.

Solubility in water — ionic substances tend to be soluble in water but insoluble in organic solvents (organic solvents include ethanol/alcohol and hydrocarbons such as hexane and those found in petrol).
Reason (extension): water is a covalent molecule, but the electrons in its bonds are more attracted towards the oxygen end of the bond, making the oxygen slightly negative and the hydrogen slightly positive — the molecule is called polar. This means reasonably strong forces can form between water molecules and ions, providing the energy to break the lattice apart.
Not all ionic substances are soluble: magnesium oxide is not soluble in water because the attractions between water molecules and the ions aren't strong enough to overcome the very strong electrostatic attractions between magnesium and oxide ions.
Hexane is non-polar and does not form strong enough attractions to the ions to break apart the ionic lattice.

The electrical conductivity of ionic substances

Ionic compounds do not conduct electricity when solid because the ions are fixed in position and are not free to move around.
They do conduct electricity when molten or dissolved in water (in aqueous solution), because the ions then become free to move around.
Molten — just means that the salt has been melted; it is a liquid.

💡 Exam tip: Use the correct words when explaining conductivity — do not use the word 'electrons'. You must talk about the ions being free to move.

✅ Quick check

Explain what is meant by (i) an ion and (ii) ionic bonding.
Draw a dot-and-cross diagram (outer electrons only) to show the ions formed when magnesium combines with fluorine (magnesium fluoride).
Work out the formula of: (a) iron(III) hydroxide, (b) ammonium sulfate, (c) aluminium oxide.
Explain why sodium chloride (a) has a high melting point, (b) does not conduct electricity when solid, (c) conducts electricity when molten.
Which has the higher melting point — potassium chloride or calcium oxide? Give a reason.

Answers

1. (i) An ion is a charged particle formed when atoms (or groups of atoms) lose or gain electrons. (ii) Ionic bonding is the strong electrostatic attraction between positive and negative ions. 2. The Mg atom [2,8,2] loses its 2 outer electrons (one to each F atom [2,7]), forming Mg²⁺ and two F⁻ ions, each ion now with a noble gas configuration — formula MgF₂. Draw two F⁻ around one Mg²⁺, showing the transfer of the two outer electrons. 3. (a) Fe³⁺ + OH⁻ → **Fe(OH)₃** (brackets needed — 3 hydroxide ions per Fe³⁺); (b) 2 × NH₄⁺ + SO₄²⁻ → **(NH₄)₂SO₄**; (c) Al³⁺ + O²⁻ → **Al₂O₃**. 4. (a) Strong electrostatic forces of attraction between oppositely charged ions in the giant ionic lattice require a lot of energy to break. (b) When solid, the ions are fixed in position and cannot move. (c) When molten the lattice breaks down and the ions become free to move and carry charge. 5. **Calcium oxide** — it contains 2+ and 2− ions, so the electrostatic attractions holding the lattice together are much stronger than the 1+ / 1− attractions in potassium chloride, so more energy is needed to break them.

Ch 8 Covalent Bonding

Chapter 8: Covalent Bonding

There are a lot more covalent compounds than ionic compounds, so it is important to understand how this bonding works. In a covalent bond, atoms are held together by sharing a pair of electrons (rather than transferring them, as in ionic bonding). This chapter covers dot-and-cross diagrams for many molecules, the difference between simple molecular structures and giant covalent structures, and why the two types behave so differently (melting/boiling points, conductivity, hardness) — including diamond, graphite and C₆₀ fullerene. Water is a covalent compound, but the salt dissolved in sea water is an ionic compound.

What is a covalent bond?

In any bond, particles are held together by electrostatic attractions between something positively charged and something negatively charged.
In a covalent bond, a pair of electrons is shared between two atoms.
What holds the atoms together is the strong electrostatic attraction between the nuclei (positively charged) of the atoms that make up the bond, and the shared pair of electrons (negatively charged).

📊 Key diagram (Fig 8.2): The shared pair of electrons — the shared electron pair sits between the two atoms and is attracted to the nuclei of both atoms: nucleus A is attracted to the electron pair, and nucleus B is also attracted to the electron pair. (In most simple examples each atom supplies one electron to the shared pair; that doesn't have to be the case — both electrons may come from the same atom.)

Covalent bonding in a hydrogen molecule

Hydrogen atoms form diatomic molecules with the formula H₂ — the two atoms in an H₂ molecule are joined by one covalent bond, and the covalent bond between two hydrogen atoms is very strong.
Diatomic — describes molecules containing two atoms (e.g. hydrogen molecules are diatomic).
Molecule — a fixed number of atoms joined together by covalent bonds. Hydrogen molecules contain two atoms; other sorts of molecule may have as many as thousands of atoms joined together (for example proteins and DNA).
Covalent bonds are often shown using dot-and-cross diagrams: although the electrons are drawn as dots or crosses, there is absolutely no difference between them in reality — the dots and crosses simply show the electrons have come from two different atoms.

📊 Key diagram (Fig 8.3): Dot-and-cross diagram for H₂ — two H atoms, each with its single electron (one drawn as a dot, one as a cross), sharing the pair so each H has 2 electrons in its outer shell.

The significance of noble gas structures in covalent bonding

In H₂, each hydrogen atom has only one electron to share, so it can only form one covalent bond. The shared pair of electrons is in the outer shell of both atoms, so each atom has the same number of electrons as a noble gas atom (helium in this case).
Does that mean the hydrogen has turned into helium? No. The number of protons in the nucleus hasn't changed — and it is the number of protons that defines what an atom is.
In virtually all molecules at this level, electrons are shared so that H atoms have a total of 2 electrons in their outer shell and all other atoms have 8 electrons in their outer shell.
Octet rule — the idea that atoms share electrons so that they end up with 8 electrons in their outer shell (some people talk about this '8'). Remember: we are counting shared electrons as belonging to the outer shells of both atoms.
When there is one atom in the middle with other atoms joined to it (as in CH₄ or PCl₃), the outer atoms will virtually always have 8 electrons in their outer shell (or 2 if they are H). It is very difficult to think of an example where the outer atoms do not have 8 electrons.
There are some molecules where the central atom does not have 8 electrons in its outer shell — we will look at a couple of examples later (BF₃).
Atoms with 8 electrons in their outer shell have the same number of electrons as a noble gas atom — they are isoelectronic with a noble gas atom.

Why does hydrogen form molecules?

In chemistry, 'stable' refers to energy: generally, the lower the energy something has, the more stable it is.
Whenever a bond is formed (of whatever kind), energy is released, and that makes the things involved more stable than they were before. The more bonds an atom can form, the more energy is released and the more stable the system becomes.
The H₂ molecule is much more stable than two separate hydrogen atoms. When bonds form, energy is given out, so the substance formed has less energy — therefore it is more stable.
Chemical reactions usually occur so that something becomes more stable. (Analogy: hold a book and let go — it falls to the floor, where it has less potential energy.)

Covalent bonding in a hydrogen chloride molecule

A chlorine atom has 7 electrons in its outer shell. By sharing 1 electron with a hydrogen atom, both atoms end up with the same number of electrons as the nearest noble gas atom:
The chlorine's arrangement around it is now 2, 8, 8 — the same as an argon atom.
The hydrogen now has 2 electrons in its outer shell — the same as helium.
Notice that only the outer-shell electrons of the chlorine are used in bonding. At International GCSE, inner electrons never get used in bonding and are often left out of bonding diagrams.

💡 Exam tip: Only leave out the inner electrons if the question tells you to. Also be careful: a line used to represent a covalent bond (as in H–Cl) is not a dot-and-cross diagram.

📊 Key diagrams: Fig 8.4 — dot-and-cross diagram for HCl (H sharing its electron with one of Cl's 7 outer electrons); Fig 8.5 — the bonding in HCl shown as H:Cl: (outer-shell electrons only); Fig 8.6 — the line between the atoms (H–Cl) represents a covalent bond.

Covalent bonding in a chlorine molecule (Cl₂)

A chlorine atom has 7 electrons in its outer shell. Each Cl shares 1 electron so that both Cl atoms have 8 electrons in their outer shell.

📊 Key diagram (Fig 8.7): Two ways of showing the bonding in Cl₂ — (a) dot-and-cross diagram of the two Cl atoms sharing a pair of outer electrons; (b) the line diagram Cl–Cl where the line represents the shared pair.

Covalent bonding in methane, ammonia and water

📊 Key diagram (Fig 8.8): Electronic configurations of C, N and O atoms — carbon [2,4], nitrogen [2,5], oxygen [2,6].

Methane (CH₄)

A carbon atom has 4 electrons in its outer shell. By sharing 1 electron with each of 4 hydrogen atoms, the C ends up with 8 electrons in its outer shell and each H ends with 2. Therefore C forms 4 covalent bonds — 1 with each H atom. Methane has the formula CH₄.

📊 Key diagram (Fig 8.9): Dot-and-cross diagrams for methane — central C sharing one pair with each of four surrounding H atoms (drawn as H–C–H in a cross arrangement).

Ammonia (NH₃)

A nitrogen atom has 5 electrons in its outer shell. It shares 3 other electrons to have 8 in its outer shell, forming 3 covalent bonds — 1 with each H atom. The formula of ammonia is NH₃.

📊 Key diagram (Fig 8.10): Dot-and-cross diagram for ammonia — N sharing one pair with each of three H atoms (plus one lone pair of N's own electrons not involved in bonding).

Water (H₂O)

An oxygen atom has 6 electrons in its outer shell. It shares 2 other electrons to have 8 in its outer shell, forming 2 covalent bonds — 1 with each H atom. The formula of water is H₂O.

📊 Key diagram (Fig 8.11): Dot-and-cross diagram for water — O sharing one pair with each of two H atoms, leaving two lone pairs.

Shapes of molecules (Looking ahead)

By understanding the bonding in a covalent molecule it is possible to work out its shape: pairs of electrons in the outer shell of the central atom repel each other and tend to get as far away from each other as possible.
In a methane molecule there are four pairs of electrons around the central C atom; to be as far apart as possible they must be arranged in a tetrahedral shape.
Tetrahedron (adjective: tetrahedral) — a triangular pyramid.
There are also four pairs of electrons around the central atom in water: two are bonding pairs and two are lone pairs of electrons (pairs not involved in bonding). All four pairs are arranged tetrahedrally, but because only two of them hold atoms, the actual shape of a water molecule is described as 'bent' or 'V-shaped'.
Because the electrons in the bonds are attracted to a different extent by the oxygen and hydrogen atoms, a water molecule is polar — it has a slightly negative end and a slightly positive end.

📊 Key diagrams: Fig 8.12 — (a) a methane molecule is tetrahedral, (b) a water molecule is bent; Fig 8.13 — a stream of water being bent by an electrically charged comb (evidence that water is polar).

Covalent bonding in ethane (a slightly more complicated molecule)

Ethane has the formula C₂H₆. The bonding is similar to methane, except that there is a carbon–carbon covalent bond as well as the carbon–hydrogen bonds.
This is an organic compound (you learn more about molecules like this in Unit 4).

📊 Key diagram (Fig 8.14): Dot-and-cross diagram for ethane — two C atoms joined by a C–C bond, each C also bonded to three H atoms (each H sharing one electron with C).

💡 Exam tip: When drawing molecules containing carbon and hydrogen it is useful to remember that carbon always forms 4 bonds (sharing 4 electrons), hydrogen always forms 1 bond (sharing 1 electron), and the hydrogen atoms always go on the outside, never in the middle.

Multiple covalent bonding

Double bonding in an oxygen molecule (O₂)

An oxygen atom has 6 electrons in its outer shell. If two oxygen atoms combine, they will both share 2 electrons each, so each atom will have 8 electrons in its outer shell.
There are therefore two shared pairs of electrons between the oxygen atoms — called a double covalent bond, or usually just a double bond.

📊 Key diagram (Fig 8.15): (a) dot-and-cross diagram for O₂ with two shared pairs between the O atoms; (b) the double covalent bond drawn as O=O (two lines).

The triple bond in a nitrogen molecule (N₂)

A nitrogen atom has 5 electrons in its outer shell. If two nitrogen atoms combine they will both share 3 electrons each, so each atom has 8 electrons in its outer shell.
There are therefore three shared pairs of electrons between the nitrogen atoms — called a triple covalent bond or usually just a triple bond.
Nitrogen gas consists of nitrogen molecules bonded like this. The triple bond is very strong and needs a lot of energy to break — that is why nitrogen is relatively unreactive.

📊 Key diagram (Fig 8.16): N₂ has a triple bond — drawn N≡N (three lines between the N atoms).

Covalent double bonding in carbon dioxide (CO₂)

An oxygen atom has 6 electrons in its outer shell and a carbon atom has 4. Each oxygen atom will share 2 electrons with the carbon atom.
Two double bonds are formed between the carbon and the two oxygens; all atoms have 8 electrons in their outer shells.

📊 Key diagram (Fig 8.17): CO₂ has two double bonds — drawn O=C=O.

The double bond in ethene (C₂H₄)

Ethene is rather like ethane, except that it only has two hydrogen atoms attached to each carbon atom, and a double bond between the carbon atoms.

📊 Key diagram (Fig 8.18): Ethene has a double bond between the C atoms — drawn H₂C=CH₂ (each C bonded to 2 H atoms and joined to the other C by a double bond).

💡 Exam tip: With organic compounds such as ethane and ethene you have to look at their names very carefully — even one different letter in the name can matter. Here, ethane and ethene are completely different compounds (you find out more in Unit 4).

Organic molecules containing halogen atoms

Bromomethane has the formula CH₃Br: the 3 H atoms and the Br atom are joined to the central C atom.
Br has 35 electrons and we haven't learned how to work out the electronic configuration of an atom with 35 electrons — but looking at the Periodic Table, Br is in Group 7, so it has 7 electrons in its outer shell and will share 1 electron so it has 8 in its outer shell. Therefore Br forms just 1 covalent bond.
When drawing a dot-and-cross diagram for CH₃Br you will only be asked to show the outer electrons.

📊 Key diagram (Fig 8.19): Dot-and-cross diagram for CH₃Br showing outer shells only — central C sharing one pair with each of 3 H atoms and one pair with the Br atom (you do not need to use different colours in the exam, even though different colours may be used in the book for clarity).

Probably the most complicated molecule you could be asked to draw a dot-and-cross diagram for is something like chloroethene (CH₂CHCl). When drawing it, remember: C forms 4 covalent bonds, H forms 1 and Cl forms 1.
There is a double bond between the 2 C atoms. It does not matter where you put the H and Cl atoms relative to the C, and you don't have to use different colours.
The structure of chloroethene can also be shown with lines representing the covalent bonds (Fig 8.21).

⚠️ Exam warning: A line-bond structure (Fig 8.21) is not a dot-and-cross diagram — you must not draw this if you are asked to draw a dot-and-cross diagram.

Some more difficult molecules where the central atom does not have 8 electrons in its outer shell

Although the outer atoms in a molecule will always have 8 (or 2 in the case of hydrogen) electrons in their outer shell, there are a few examples where the central atom has more or fewer than 8 electrons in its outer shell.

Boron trifluoride (BF₃)

Boron is the central atom; the F atoms are the outer atoms. Each F shares 1 electron so it has 8 in its outer shell, but this means B only has a total of 6 electrons in its outer shell.
Another way of thinking about this: a B atom only has 3 electrons in its outer shell, and so this is the maximum number that it can share.

📊 Key diagram (Fig 8.22): Dot-and-cross diagram for BF₃ — three F atoms each sharing a pair with the central B, which ends with only 6 outer electrons (it is electron-deficient).

Sulfur dioxide (SO₂)

Also called sulfur(IV) oxide. The central atom is S and the outer atoms are O.
Each O atom has 6 electrons in its outer shell and needs to share 2 more to have 8; so the S atom shares 2 electrons with each O atom to form two double bonds.
A sulfur atom originally had 6 electrons in its outer shell; after sharing 4 electrons it has 10 electrons in its outer shell.

📊 Key diagram (Fig 8.23): There are two double bonds in SO₂ (O=S=O).

Atoms in Periods 3 and below (Periods 4, 5, 6, 7) can have more than 8 electrons in their outer shells. The maximum number of electrons the central atom in a molecule can share is equal to the number of electrons in its outer shell — so sulfur can form up to 6 bonds (in SF₆) and chlorine up to 7 (in HClO₄).

Simple molecular structures

Molecule — a fixed number of atoms joined by strong covalent bonds. If we look closely at liquid water, there are individual water molecules where the H and O atoms are joined by strong covalent bonds — but there must also be some forces between water molecules which keep them in the liquid state.
Intermolecular forces — the forces of attraction between molecules ('intermolecular' literally means 'between molecules'). These are much weaker than covalent bonds.
When we boil water, only these weak intermolecular forces of attraction are broken; covalent bonds are not broken. The covalent bonds between the H and O atoms in the molecules have not changed in any way — all that has changed in gaseous water is that the intermolecular forces have been broken.
Simple molecular structure — when a substance consists of molecules with (weak) intermolecular forces of attraction between them.
Virtually all compounds that have covalent bonding have simple molecular structures.
Examples of substances with simple molecular structures: H₂O, CO₂, CH₄, N₂ and C₂H₄ — all the things we have drawn dot-and-cross diagrams for above.

📊 Key diagram (Figs 8.24 & 8.25): Water as a simple molecular compound — strong covalent bonds join the H and O atoms within each molecule, but the intermolecular attractions between water molecules are weaker. On boiling/evaporating, only the intermolecular forces are broken — H₂O molecules exist unchanged in both liquid and gaseous water.

Substances with simple molecular structures tend to be gases or liquids, or solids with low melting and boiling points. The reason: not much energy is required to break the weak intermolecular forces of attraction. Remember — no covalent bonds are broken; covalent bonds are strong.

Melting and boiling points increase as relative molecular mass increases

The halogens (Group 7) all have a simple molecular structure consisting of diatomic molecules with intermolecular forces between them.
Their melting and boiling points increase as the relative molecular mass increases:

Halogen	Formula	Relative molecular mass (Mᵣ)	Melting point/°C	Boiling point/°C
fluorine	F₂	38	−220	−188
chlorine	Cl₂	71	−101	−34
bromine	Br₂	160	−7	59
iodine	I₂	254	114	184

Explanation: as we melt or boil these substances we are only breaking the intermolecular forces of attraction between molecules. As the Mᵣ increases down the group, more energy is needed to break the intermolecular forces, which means the intermolecular forces of attraction become stronger as the relative molecular mass increases.
This is seen quite often: boiling points increase along the series CH₄, C₂H₆, C₃H₈ as relative molecular mass increases (also true of the alkanes, Chapter 23).
⚠️ It is not always the case that melting and boiling points increase as Mᵣ increases — the rule only really applies to sets of very similar substances, such as the halogens or the alkanes. Examples where the rule does not work: water (Mᵣ = 18, boiling point 100 °C), ethane (Mᵣ = 30, b.p. −89 °C), ammonia (Mᵣ = 17, b.p. −33 °C) and PH₃ (Mᵣ = 34, b.p. −88 °C).
There is a special type of intermolecular force between water molecules called hydrogen bonding — it gives water some of its very special properties, for example the solid form (ice) is less dense than the liquid form (discussed further in Chapter 23).

Some other physical properties of covalent compounds

Covalent molecular compounds do not conduct electricity. This is because the molecules don't have any overall electrical charge (there are no ions) and all the electrons are held tightly in the atoms or in covalent bonds, so they are not able to move from molecule to molecule.
Covalent molecular substances tend to be insoluble in water. There are some exceptions, for example ethanol (C₂H₅OH) and substances such as NH₃ and HCl that react with water as they dissolve.
Covalent molecular substances are often soluble in organic solvents.

Giant covalent structures

Not all covalent substances have simple molecular structures — some form giant covalent structures that continue on and on in three dimensions.

Diamond

Diamond is a form of pure carbon.
Each carbon atom has four electrons in its outer shell, so it forms four covalent bonds — bonding strongly to four other carbon atoms in a tetrahedral arrangement.
Tetrahedron — a triangular-based pyramid: one atom at the centre of the tetrahedron, with the atoms it is attached to at the four corners.
A real diamond crystal is a giant covalent structure; it continues on and on in three dimensions, and every atom drawn is attached to four others. It is not a molecule because the number of atoms joined up is completely variable and depends on the size of the crystal, whereas molecules always contain fixed numbers of atoms joined by covalent bonds.

📊 Key diagram (Fig 8.26): Part of the structure of diamond — each C atom bonded to four others in a tetrahedral arrangement (some atoms in the diagram may appear to form only one or two bonds, but that's only because we are showing a small part of the whole 3-D structure). Each line represents a covalent bond.

Very high melting and boiling points — because of the very strong carbon–carbon covalent bonds extending throughout the whole crystal in three dimensions. A lot of energy has to be supplied to break these strong covalent bonds.
How this differs from simple molecular structures: to melt or boil a substance with a simple molecular structure (e.g. CH₄) we only had to supply enough energy to break the relatively weak intermolecular forces. In diamond there are no intermolecular forces (it has a giant structure — there are no molecules); the very strong covalent bonds themselves must be broken in order to melt or boil it.
In general, all substances with giant covalent structures are solids with high melting and boiling points, because a lot of energy has to be supplied to break all the strong covalent bonds throughout the giant structure.
Other substances with giant covalent structures include graphite (discussed below) and silicon dioxide (SiO₂).
Diamond is very hard — again, a lot of energy has to be supplied to break the strong covalent bonds. Drill bits can be tipped with diamonds for drills used on stone and rock. Diamond is crystalline and is the hardest naturally occurring substance.
Diamond does not conduct electricity — all the electrons in the outer shells of the carbon atoms are tightly held in covalent bonds between the atoms; none are free to move around.
Diamond doesn't dissolve in water or any other solvent — because of the strong covalent bonds between the carbon atoms; if the diamond dissolved, these bonds would have to be broken.
Diamond conducts heat very well (better than any other element): as one end of the crystal is heated the atoms vibrate more, and the strong bonds throughout the giant structure mean these vibrations are quickly transmitted from one end of the crystal to the other.

Graphite

Graphite is also a form of carbon, but the atoms are arranged differently, although it still has a giant structure.
Graphite has a layer structure, rather like a pack of playing cards — in a pack, each card is strong but the individual cards are easily separated; the same is true in graphite.
Within each layer, each carbon atom is joined to only three others, forming strong covalent bonds. The attractions between the layers are much weaker — not much energy is needed to overcome them.
Graphite is soft — although the forces holding the atoms together in each layer are very strong, the attractions between layers are weak, so the layers slide over each other and can easily be flaked off.
Uses of this softness: graphite (mixed with clay to make it harder) is used in pencils — when you write you leave a trail of graphite layers behind on the paper. Pure graphite is so slippery it is used as a dry lubricant, e.g. powdered graphite lubricates locks.
Graphite has high melting and boiling points — to melt/boil graphite you don't just have to separate the layers; you have to break up the whole structure, including the strong covalent bonds, and that needs very large amounts of energy.
Graphite conducts electricity. Each carbon atom uses three of its outer-shell electrons to form three single covalent bonds; the fourth electron in the outer shell of each atom is free to move around throughout the whole of the layer. These free-moving electrons are called delocalised electrons, and their movement allows graphite to conduct electricity. (Delocalised electrons also explain why graphite electrodes in an electric arc furnace used to produce steel glow red hot.)
Graphite is insoluble in all solvents, because it would take too much energy to break all the strong covalent bonds.
Graphite is less dense than diamond because the layers in graphite are relatively far apart — the distance between the graphite layers is more than twice the distance between atoms within each layer. In a sense a graphite crystal contains a lot of wasted space, which isn't there in a diamond crystal.

📊 Key diagram (Figs 8.28–8.30): The structure of graphite — (a) layered structure like a pack of playing cards; (b) within each layer, each carbon atom joins to three others, with big gaps between layers; (c) edge-on view: the forces of attraction between layers are drawn as dashed lines — these are not bonds and could have been drawn anywhere between the layers.

💡 Exam tip: The reason graphite layers slide over each other fairly easily is actually more complicated — graphite is not a lubricant in a vacuum; its lubricating action relies on water molecules sticking to the surface, which does not happen in a vacuum.

C₆₀ fullerene

Allotropes — different forms of the same element. Diamond and graphite are two allotropes of carbon; another allotrope is C₆₀ fullerene.
Diamond and graphite both have giant structures, but C₆₀ fullerene has a simple molecular structure, which has a big influence on its physical properties.
In solid or liquid C₆₀ fullerene there are C₆₀ molecules with weak intermolecular forces between them. Each C₆₀ fullerene molecule contains 60 carbon atoms joined by covalent bonds (a fixed number of atoms — it is a molecule).
There are different fullerenes containing different numbers of carbon atoms — this is why we include the C₆₀ in the name.
C₆₀ fullerene has lower melting and boiling points than diamond and graphite. When fullerene melts, only the relatively weak intermolecular forces of attraction must be broken — this does not require as much energy as breaking all the strong covalent bonds when diamond and graphite melt.
C₆₀ fullerene is not as hard as diamond — it does not take as much energy to break the intermolecular forces of attraction in C₆₀ compared with breaking the strong covalent bonds in diamond.
C₆₀ fullerene does not conduct electricity — although all the carbon atoms in C₆₀ only form three bonds, the fourth electron on each atom can only move around within each C₆₀ molecule; the electrons cannot jump from molecule to molecule.
Unlike diamond and graphite, C₆₀ fullerene does dissolve in some solvents — only relatively weak intermolecular forces of attraction have to be broken for it to dissolve.

📊 Key diagrams: Fig 8.32 — a C₆₀ fullerene molecule (60 carbon atoms arranged in a cage of hexagons and pentagons, like a football); Fig 8.33 — C₆₀ fullerene has a simple molecular structure: separate C₆₀ molecules with intermolecular forces between them.

✅ Quick check

State whether each of the following compounds is ionic or covalent: (a) MgO, (b) CH₃Br, (c) H₂O₂, (d) FeCl₃, (e) NaF, (f) HCN.
Explain what is meant by a covalent bond and how it holds two atoms together.
Explain why carbon dioxide sublimes at −78.5 °C but diamond sublimes at around 4000 °C.
Explain in terms of structure and bonding why: (a) diamond is harder than graphite, (b) graphite conducts electricity but diamond does not.
The boiling points of the halogens rise from fluorine (−188 °C) to iodine (184 °C). Explain why.
Why does C₆₀ fullerene have a much lower melting point than graphite?

Answers

1. Ionic: MgO, FeCl₃, NaF. Covalent: CH₃Br, H₂O₂, HCN (all non-metal compounds). 2. A covalent bond is a shared pair of electrons between two atoms. It holds the atoms together because of the strong electrostatic attraction between the positively charged nuclei of both atoms and the negatively charged shared pair of electrons. 3. CO₂ has a simple molecular structure — only weak intermolecular forces between molecules need breaking, so little energy is needed. Diamond is a giant covalent structure: to sublime it you must break the very strong covalent bonds between carbon atoms throughout the whole 3-D structure, which needs a huge amount of energy. 4. (a) In diamond every carbon is bonded to four others in a rigid tetrahedral network, so a large force is needed to break strong C–C bonds; in graphite each carbon bonds to only three others in layers held by weak forces between layers, so layers slide and flake easily. (b) In graphite, each carbon has a fourth outer electron that is delocalised and free to move throughout each layer, so it conducts; in diamond all four outer electrons are held tightly in covalent bonds, so none are free to move. 5. All the halogens are simple molecular substances (diatomic molecules). As the relative molecular mass increases down the group, the intermolecular forces of attraction between molecules become stronger, so more energy is needed to break them when boiling. 6. C₆₀ fullerene has a simple molecular structure — only weak intermolecular forces between the molecules need to be broken when it melts. Graphite is a giant covalent structure, so melting requires breaking strong covalent bonds throughout the whole structure, needing far more energy.

Ch 9 Metallic Bonding

Chapter 9: Metallic Bonding

Metals are built from a regular lattice of positive ions surrounded by a free-moving 'sea' of delocalised electrons. The electrostatic attraction between these two holds the metal together, and explains why metals are strong, conduct electricity, and can be hammered into shape or drawn into wires. (Whole chapter is CHEMISTRY ONLY.)

Learning objectives: know how to represent a metallic lattice as a 2D diagram • understand metallic bonding in terms of electrostatic attractions • explain typical physical properties of metals, including electrical conductivity and malleability.

Metallic bonding

Metal structure — when metal atoms (e.g. sodium) bond together to form the solid, each atom's outer electron becomes free to move throughout the whole structure.
Delocalised electrons — electrons that are no longer attached to particular atoms or pairs of atoms; you can think of them as flowing around through the whole metal.
When a sodium atom loses its outer electron, a sodium ion (Na⁺) is left behind.
A metallic structure therefore consists of a lattice (regular arrangement) of positive ions in a sea of delocalised electrons (Figure 9.3).
Metallic bonding — the electrostatic forces of attraction between each positive ion and the delocalised electrons; this is what holds the structure together (Figure 9.4).
Metals have giant structures: there are no individual molecules, and all the positive ions in the lattice attract all the delocalised electrons.
The charge on the metal ion depends on the number of electrons the original atom had in its outer shell: Group 1 elements form 1+ ions, Group 2 elements form 2+ ions.
Thinking about the structure as a whole, the number of electrons exactly balances the number of positive charges, so the metal as a whole carries no charge.

📊 Key diagram: metallic lattice (2D)

        sea of delocalised electrons (e⁻ moving freely)
    e⁻      e⁻      e⁻      e⁻
      Na⁺     Na⁺     Na⁺          ← lattice of positive
    e⁻      e⁻      e⁻      e⁻        metal ions in a
      Na⁺     Na⁺     Na⁺             regular arrangement
    e⁻      e⁻      e⁻      e⁻

  Na⁺ ┄┄┄┄ e⁻  ←  electrostatic attraction between the
       ┄┄┄┄┄┄      positive ions and negative electrons

Different metals have different arrangements of ions in the lattice — but for a diagram you only need to draw the ions in a regular arrangement with delocalised electrons between them; that will be fine.

💡 Exam tip (from the book's Hints): - When drawing a metallic structure, don't worry about the exact arrangement of ions — a regular arrangement of positive ions in a sea of electrons is enough. - In equations, write a metal such as sodium as atoms (Na), not as Na⁺ or Na⁻ — students who know about metallic bonding sometimes worry about which to write; you write it as atoms, e.g. Na. - When explaining bonding strength, compare the charge on the ions and the number of electrons delocalised per atom.

Physical properties of metals

Most metals are hard and have high melting points — this suggests the electrostatic forces of attraction between the positive ions and the delocalised electrons are strong.
In sodium, only one electron per atom is delocalised, leaving ions with only a single positive charge. The bonding is therefore quite weak, which is why sodium is fairly soft, with a low melting point for a metal.
In magnesium, both outer electrons are delocalised into the 'sea', leaving behind 2+ ions. There is a much stronger electrostatic attraction between the 2+ ions and the delocalised electrons, so the bonding is stronger and magnesium's melting point is higher.

💡 Exam tip: Na → Mg → Al trend: sodium, magnesium and aluminium have boiling points of 890 °C, 1110 °C and 2470 °C respectively (from the chapter questions). Explain the increase using electronic configurations (1, 2 and 3 outer electrons → 1+, 2+ and 3+ ions → more electrons delocalised per atom → stronger electrostatic attraction between ions and delocalised electrons → more energy needed to break the metallic bonds).

Metals conduct electricity

Metals conduct electricity because the delocalised electrons are free to move throughout the structure.
When a piece of metal is attached to a power source: electrons are attracted to the positive terminal and flow away along the wire; more electrons flow in from the negative terminal of the power source to replace those moving away (Figure 9.5).
Because electrons are flowing in the external circuit, an electric current flows.
KEY POINT — don't confuse metals with ionic substances:
Metals conduct when solid or molten because the delocalised electrons are free to move.
Ionic substances conduct only when molten or in aqueous solution because then the ions are free to move; ionic substances do not conduct when solid because the ions are not free to move.
In both cases charged particles must be free to move for conduction to occur.

Metals are malleable

Malleable — can be hammered into different shapes (metals can be hammered into shape).
When a force is applied to a piece of metal, the layers of positive ions slide over each other.
This does not affect the bonding in the structure — the positive ions are still attracted to the delocalised electrons, so the structure does not break (Figure 9.6).
Ductile — can be drawn out into wires. The explanation is the same as for malleability: layers of ions slide over one another without breaking the bonding.

📊 Key diagram: why metals are malleable

  layers of positive ions before force          after force (layers slide)
   Na⁺  Na⁺  Na⁺  Na⁺      e⁻ sea               Na⁺  Na⁺  Na⁺  Na⁺
   Na⁺  Na⁺  Na⁺  Na⁺  ┄┄┄┄┄┄┄┄┄┄┄┄            ── Na⁺  Na⁺  Na⁺ ──  ← ions
   Na⁺  Na⁺  Na⁺  Na⁺                              slide over each
   Na⁺  Na⁺  Na⁺  Na⁺      ← force applied        other; bonding to
                                                    e⁻ sea unchanged

💡 Exam tip: to answer "explain why all these metals are malleable/conduct": state the structure (lattice of positive ions in a sea of delocalised electrons) and then link it to the property (layers can slide over each other without breaking the bonding / the delocalised electrons are free to move and carry charge).

✅ Quick check

Q: Describe, with a 2D diagram, the structure of a metal such as sodium. A: A regular lattice of positive ions (Na⁺) surrounded by a sea of delocalised electrons (the outer electron of each atom is free to move throughout the whole structure).
Q: Define metallic bonding. A: The electrostatic forces of attraction between the positive metal ions and the delocalised electrons; this holds the metallic structure together.
Q: Magnesium has a higher melting point than sodium. Explain why. A: Magnesium has two outer electrons, both delocalised, leaving 2+ ions; sodium has one, leaving 1+ ions. The stronger electrostatic attraction between the 2+ ions and the delocalised electrons in magnesium means stronger metallic bonding, so more energy is needed to break it.
Q: Why do metals conduct electricity when solid, but ionic compounds only when molten or dissolved? A: In metals the delocalised electrons are free to move even in the solid. In ionic compounds the ions are locked in the lattice when solid and only become free to move when molten or in aqueous solution.
Q: What does 'malleable' mean and why are metals malleable? A: Malleable means the metal can be hammered into different shapes; the layers of positive ions slide over each other when a force is applied, without breaking the bonding, because the ions are still attracted to the delocalised electrons.

Ch 10 Electrolysis

Chapter 10: Electrolysis

Passing electricity through an ionic compound that is molten or in solution causes a chemical change called electrolysis: positive ions (cations) are attracted to the cathode and negative ions (anions) to the anode, where they are discharged. This chapter gives the rules for predicting products — for molten compounds you get the metal and non-metal, but in aqueous solution the ions from water (H⁺ and OH⁻) compete, so hydrogen is often produced at the cathode and oxygen at the anode. (Whole chapter is CHEMISTRY ONLY.)

Learning objectives: understand why ionic compounds conduct electricity only when molten or in aqueous solution, and why covalent compounds do not conduct electricity • know that anion and cation mean negative and positive ions • write ionic half-equations for the reactions at the electrodes and classify them as oxidation or reduction • describe experiments to investigate electrolysis (molten lead(II) bromide; aqueous sodium chloride, dilute sulfuric acid and copper(II) sulfate) using inert electrodes, and predict the products • practical: investigate the electrolysis of aqueous solutions.

Why things conduct electricity

For anything to conduct electricity there must be charged particles present AND those charged particles must be free to move.
The charged particles are either electrons or ions — be clear which one you are talking about.

Metals

Structure: a lattice of positive ions surrounded by a sea of delocalised electrons.
Metals conduct because the delocalised electrons are free to move.

Ionic compounds

Ionic compounds (e.g. sodium chloride, potassium iodide) do not conduct when solid: the ions are held tightly in the lattice — they cannot move around, only vibrate.
They do conduct when molten (melted) or in aqueous solution (dissolved in water), because the ions become free to move.
Terminology:
Cations — positive ions (attracted to the cathode).
Anions — negative ions (attracted to the anode).

Covalent compounds

Covalent molecular compounds (e.g. hexane C₆H₁₄, methane CH₄, carbon dioxide CO₂) do not conduct electricity in any state or in solution:
the molecules have no overall electrical charge, so there are no charged particles to move;
all the electrons are held tightly in atoms or in covalent bonds and cannot move from molecule to molecule.
Exceptions — covalent compounds that form ions as they react with water, e.g.:
Ammonia: NH₃(g) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq) — ammonia solution conducts because there are ions free to move. (⇌ = reversible reaction — see Chapter 21.)
Hydrogen chloride gas dissolves in water to form hydrochloric acid: HCl(aq) → H⁺(aq) + Cl⁻(aq) (HCl ionises in water).

Passing electricity through compounds: electrolysis

When metals conduct electricity nothing chemical happens — the metal just gets hot.
When you pass electricity through an ionic compound that is molten or in solution, a chemical reaction occurs (e.g. a solution of potassium iodide gives hydrogen gas and iodine).
Electrolysis — a chemical change caused by passing an electric current through a compound which is either molten or in solution.

Some other important words

Electrolyte — a liquid or solution that undergoes electrolysis. Electrolytes all contain ions; the movement of the ions is responsible for both the conduction of electricity and the chemical changes that take place.
Electrodes — the two conductors through which electricity is passed into and out of the electrolyte. Carbon is frequently used because it conducts electricity and is chemically fairly inert (does not react with things); platinum is also fairly inert. Various other metals are sometimes used as well.
Anode — the positive electrode.
Cathode — the negative electrode.

💡 Exam tip / mnemonics: - PANIC: Positive Anode, Negative Is Cathode. - AN OX, RED CAT: ANions → OXidation at the anode; REDuction at the CAThode. - Reduction always occurs at the cathode; oxidation always occurs at the anode.

The electrolysis of molten compounds

🔬 Practical: Electrolysing molten lead(II) bromide, PbBr₂

Apparatus (Figure 10.2): a crucible containing lead(II) bromide, two carbon (graphite) electrodes connected to a 6 V DC power supply (battery or power pack — the voltage isn't critical) with a bulb in the external circuit, and a heat source.

Method / observations (Figure 10.3): - Nothing at all happens until the lead(II) bromide melts. - As soon as it melts and electricity passes: the bulb lights up, showing electrons are flowing. - There is bubbling around the anode (connected to the positive terminal) — brown bromine gas is given off. - Nothing seems to happen at the cathode at first, but afterwards metallic lead is found underneath it (molten lead sinks to the bottom of the container). - When heating stops and the lead(II) bromide solidifies again, everything stops: no more bubbling and the bulb goes out.

Why: solid lead(II) bromide is a giant ionic lattice of Pb²⁺ and Br⁻ ions locked tightly in place — they can't move, so the solid doesn't conduct. Once melted, the ions are free to move and carry the current.

Explanation at the electrodes (Figures 10.4–10.5): - The power source pumps mobile electrons away from the left-hand (anode) electrode towards the right-hand (cathode) electrode, so the anode is positive (short of electrons) and the cathode is negative (extra electrons). - Cathode: Pb²⁺ ions are attracted; each lead(II) ion picks up two electrons from the electrode and forms neutral lead atoms, which fall to the bottom as molten lead:

Pb²⁺ + 2e⁻ → Pb          (reduction — gain of electrons)

Anode: Br⁻ ions are attracted; each bromide ion loses its extra electron to the electrode (the electrode is short of electrons). Two bromine atoms join in pairs to make bromine molecules:

Br⁻ → Br + e⁻
2Br → Br₂
Overall:  2Br⁻ → Br₂ + 2e⁻          (oxidation — loss of electrons)

New electrons on the anode flow back into the power source; because electrons keep flowing in the external circuit, the bulb stays lit.
Discharging an ion — an ion is discharged when it loses its charge, either by giving up electron(s) to the electrode or receiving electron(s) from it. Here, bromide ions and lead(II) ions were discharged at the electrodes.

💡 Exam tip (Key Points): - Half-equations show either oxidation or reduction — electrons are shown as e⁻. Each half-equation must balance both the atoms and the total charge on each side. - Do not confuse bromide (the ion, Br⁻) with bromine (the element/molecule, Br₂). Bromide ions are attracted to the anode, where they lose electrons to form bromine molecules. - Remember: electrons flow in the external circuit (wire, power pack, bulb, electrodes); ions flow in the electrolyte.

Electrolysis and redox

Oxidation — loss of electrons. (E.g. bromide ions are oxidised to bromine molecules.)
Reduction — gain of electrons. (E.g. lead(II) ions are reduced to lead atoms.)
Mnemonic: OILRIG — Oxidation Is Loss, Reduction Is Gain (of electrons).
If something loses electrons, something else must gain them, so oxidation and reduction always occur at the same time — these are redox reactions (REDuction + OXidation).
At the electrodes each equation shows only ONE process (oxidation or reduction), so these are called half-equations.
You may be asked to write ionic half-equations representing the electrode reactions in the exam.

📊 Key diagram: electrolysis cell (molten compound)

        DC power supply
         ┌────┤├────┐
         │   (+)(−)  │
     anode │  cathode │
     (+) ──┴──┐  ┌──┴── (−)        carbon electrodes
              │  │
         ┌────┴──┴────┐
         │  molten    │  ← electrolyte
         │  PbBr₂     │
         └────┬──┬────┘
              │  │
         bubbles of brown Br₂ at anode (+) ; molten Pb collects at bottom

Real cell: anode on the left in the book's figure is connected to +ve; bromide ions (Br⁻) move to the anode and are oxidised to Br₂ gas; Pb²⁺ ions move to the cathode and are reduced to Pb.

The electrolysis of other molten substances

Molten sodium chloride → sodium at the cathode, chlorine at the anode:

cathode:  Na⁺ + e⁻ → Na          reduction (sodium ions reduced to sodium atoms)
anode:    2Cl⁻ → Cl₂ + 2e⁻       oxidation (chloride ions oxidised to chlorine molecules)

Half-equations must balance atoms AND charge: total charge must be the same on both sides — that is why the chloride half-equation needs 2 electrons but the sodium one only 1.
Molten aluminium oxide (Al₂O₃) → aluminium at the cathode, oxygen at the anode. Aluminium is in Group 3, so an Al atom has 3 outer electrons and forms Al³⁺:

cathode:  Al³⁺ + 3e⁻ → Al        reduction
anode:    2O²⁻ → O₂ + 4e⁻        oxidation

(Each O²⁻ ion has 2 'extra' electrons and oxygen atoms go around in pairs, so 4 electrons must be removed to form O₂ — total charge 4− on each side.) - Molten zinc(II) chloride → zinc at the cathode, chlorine at the anode:

cathode:  Zn²⁺ + 2e⁻ → Zn        reduction
anode:    2Cl⁻ → Cl₂ + 2e⁻       oxidation

💡 Exam tip: a common mistake is putting the electrons on the wrong side of a half-equation. Check the charges — make sure the total charge is the same on both sides. (For example, Al³⁺ + 3e⁻ has zero total charge on the left, matching Al on the right: the three electrons are the ones the aluminium atom lost to become Al³⁺, so they must be put back to form the atom.)

Generalisation

If you electrolyse a molten ionic compound containing only two elements, you get the metal at the cathode (metals form positive ions) and the non-metal at the anode (non-metals form negative ions).
Origin of the names: cations are positive ions because they are attracted to the cathode; anions are negative ions because they are attracted to the anode.

Not all ionic compounds can be melted

Some ionic compounds break up into simpler substances before their melting point — e.g. copper(II) carbonate decomposes into copper(II) oxide and carbon dioxide even on gentle heating; it is impossible to melt it (so it cannot be electrolysed molten).

The electrolysis of aqueous solutions

When aqueous solutions are electrolysed, the products are not always the same as when the molten salt is electrolysed, because the water also supplies ions.

🔬 Practical (Activity 5): Investigating the electrolysis of aqueous solutions — concentrated sodium chloride solution

Apparatus (Figure 10.6): a glass tube with a rubber bung holding two carbon electrodes (together called an electrolytic cell), a battery/power pack, concentrated sodium chloride solution, and test-tubes to collect the gases over each electrode.

Method: 1. Pour concentrated sodium chloride solution into the glass tube. 2. Place a test-tube containing sodium chloride solution over each electrode (the test-tubes must not completely cover the electrodes, or ions will be unable to flow and there will be no current). 3. Connect the battery/power pack to the electrodes. 4. Look for bubbles of gas or a metal forming at the electrodes; test any gases.

Safety Note: wear eye protection; do the experiment in a fume cupboard or well-ventilated room because chlorine gas is poisonous; do not smell the chlorine (especially if you have asthma); once chlorine is detected the current must be switched off.

Observations: bubbles of gas at both electrodes. Testing the gases shows hydrogen forms at the cathode (negative electrode) and chlorine forms at the anode (positive electrode). The volume of chlorine collected is less than the expected equal volume of hydrogen because chlorine is more soluble in water.

The electrolysis of sodium chloride solution

Electrolysing sodium chloride solution does not give the same products as molten sodium chloride: chlorine is still formed at the anode, but hydrogen (from the water) is produced at the cathode rather than sodium.
Water is a weak electrolyte — it ionises very slightly to give hydrogen ions and hydroxide ions:

H₂O(l) ⇌ H⁺(aq) + OH⁻(aq)

(The reversible sign shows water molecules breaking up to form the ions while these ions are recombining to make water again.) - Whenever water is present you must consider these ions as well as the ions of the compound being electrolysed.

At the cathode

The solution contains Na⁺(aq) and H⁺(aq), both attracted to the negative electrode. However, sodium is a very reactive metal: it is very difficult to add an electron to a sodium ion to turn it back into a sodium atom. Hydrogen is less reactive than sodium, so it is easier to add an electron to a hydrogen ion. The hydrogen ions are discharged; each hydrogen atom formed combines with another to make a molecule:

2H⁺(aq) + 2e⁻ → H₂(g)          reduction

Each time a water molecule ionises it also produces a hydroxide ion; because the H⁺ ions are discharged (removed) as hydrogen gas, they can no longer react with the OH⁻ ions to reform water — the ionisation of water becomes a one-way process — so OH⁻ ions build up around the cathode, making the solution alkaline there.
Alternative acceptable half-equation (starting from neutral water molecules taking electrons directly from the cathode — makes the alkalinity easier to see):

2H₂O(l) + 2e⁻ → H₂(g) + 2OH⁻(aq)

Either half-equation is accepted in the exam.

At the anode

Cl⁻(aq) and OH⁻(aq) are both attracted to the positive anode. It is slightly easier to oxidise (remove electrons from) the hydroxide ion than the chloride ion, but there is little difference — and there are many more chloride ions present, so it is mainly the chloride ions that are oxidised:

2Cl⁻(aq) → Cl₂(g) + 2e⁻          oxidation

Extension work (unlikely to be asked at International GCSE): if the solution is dilute, you get noticeable amounts of oxygen produced as well as chlorine — the oxygen comes from the hydroxide ions:

4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻

The remaining solution

Ions originally present: Na⁺ and Cl⁻ (from NaCl); H⁺ and OH⁻ (from H₂O).
After long electrolysis, Cl⁻ and H⁺ have been removed by being discharged at the electrodes; we are left with Na⁺ and OH⁻ — sodium hydroxide solution (NaOH).

The electrolysis of copper(II) sulfate solution using inert electrodes

📊 Key diagram (Figure 10.8): carbon (inert) electrodes in blue copper(II) sulfate solution; cathode becomes coated with pink-brown copper; bubbles of oxygen gas at the anode.

Cathode: Cu²⁺(aq) and H⁺(aq) (from water) are both attracted. Copper is below hydrogen in the reactivity series, so it is easier to add electrons to copper ions than to hydrogen ions — the cathode becomes coated with pink-brown copper:

Cu²⁺(aq) + 2e⁻ → Cu(s)          reduction

Anode: sulfate ions and hydroxide ions (from water) are attracted. Sulfate ions aren't easy to oxidise, so instead you get oxygen from the oxidation of hydroxide ions from the water:

4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻          oxidation

Alternative acceptable half-equation (starting from water; shows why the solution becomes acidic — H⁺(aq) ions are produced):

2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻

Products: copper at the cathode and oxygen gas at the anode.
If electrolysis continues for a long time the copper(II) ions are eventually all used up, so the blue colour fades to colourless.
What is left? Original ions: Cu²⁺ and SO₄²⁻ (from CuSO₄); H⁺ and OH⁻ (from water). Copper ions and hydroxide ions are discharged; hydrogen ions and sulfate ions are not. The solution turns into dilute sulfuric acid (H₂SO₄) — electrolysis then continues as for dilute sulfuric acid (below).
Extension work: the reason inert electrodes are stressed is that what the electrode is made of can affect the products. With copper electrodes, no oxygen is produced at the anode — instead the anode gets smaller as copper ions go into solution: Cu(s) → Cu²⁺(aq) + 2e⁻. The cathode reaction is the same as with inert electrodes. (This is the basis of copper purification/electroplating.)

The electrolysis of dilute sulfuric acid using inert electrodes

Cathode: the only positive ions arriving are hydrogen ions (from the acid and the water); these are discharged to give hydrogen gas:

2H⁺(aq) + 2e⁻ → H₂(g)          reduction

Anode: sulfate ions and hydroxide ions (from the water) arrive; the sulfate ions are too difficult to oxidise, so oxygen is obtained from the oxidation of hydroxide ions from the water:

4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻          oxidation

Twice as much hydrogen as oxygen is produced. Look at the half-equations: for every 4 electrons flowing around the circuit you get 1 molecule of O₂ but 2 molecules of H₂ — twice the number of molecules occupy twice the volume.
In practice the amount of hydrogen collected is more than twice the oxygen, because oxygen is more soluble in water than hydrogen. To get better results, run the electrolysis for a few minutes first to saturate the water with oxygen, then start collecting the gases.
💡 DID YOU KNOW? This experiment (measuring the 2:1 volume ratio of H₂ to O₂) could be used to show that the formula of water is H₂O.
🔬 Apparatus (Figure 10.9): dilute sulfuric acid with platinum electrodes in a DC circuit, collecting the gases in tubes/measuring cylinders to compare volumes.

The electrolysis of some other solutions using inert electrodes — how to work out what will happen

Rule for the cathode (using the reactivity series, Figure 10.10 — most reactive at top: potassium, sodium, lithium, calcium, magnesium, aluminium, (carbon), zinc, iron, (hydrogen), copper, silver, gold):
If the metal is high in the reactivity series (more reactive than hydrogen, roughly above zinc), you get hydrogen produced at the cathode instead of the metal (because it is easier to discharge H⁺ than to turn the very reactive metal ion back into an atom).
If the metal is below hydrogen in the reactivity series (e.g. copper), you obtain the metal at the cathode.
In between (e.g. a moderately reactive metal such as zinc): reasonably concentrated solutions give you the metal; very dilute solutions give mainly hydrogen; in between you get both. (At International GCSE you probably won't have to worry about this — exam examples are always clear.)
Rule for the anode:
If you have solutions of halides (chlorides, bromides, iodides), you obtain the halogen (chlorine, bromine or iodine) at the anode.
With other common negative ions (sulfate, nitrate, hydroxide), you obtain oxygen at the anode (because the sulfate/nitrate ions are too difficult to oxidise, so hydroxide ions from the water are oxidised instead).

Table 10.1: Electrolysis of solutions using inert (carbon or platinum) electrodes

Solution	Cathode (−) product	Cathode half-equation	Anode (+) product	Anode half-equation
KI(aq)	hydrogen	`2H⁺(aq) + 2e⁻ → H₂(g)`	iodine	`2I⁻(aq) → I₂(aq) + 2e⁻`
MgBr₂(aq)	hydrogen	`2H⁺(aq) + 2e⁻ → H₂(g)`	bromine	`2Br⁻(aq) → Br₂(aq) + 2e⁻`
H₂SO₄(aq)	hydrogen	`2H⁺(aq) + 2e⁻ → H₂(g)`	oxygen	`4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻`
CuSO₄(aq)	copper	`Cu²⁺(aq) + 2e⁻ → Cu(s)`	oxygen	`4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻`

💡 Exam tip: to predict products: (1) list ALL the ions present (from the salt AND from water: H⁺, OH⁻); (2) at the cathode, H⁺ is discharged unless the metal is below hydrogen in the reactivity series (then the metal is deposited); (3) at the anode, a halide ion is discharged as the halogen, otherwise OH⁻ is discharged giving oxygen.

What would happen with non-electrolytes?

For electrolysis to work there must be ions present — the current in the external circuit can flow only if there are ions that can move and be discharged.
If you tried to electrolyse a covalent compound (molten or in solution) there would be no current flow because there aren't any ions — and nothing else would happen either.
Non-electrolyte — a substance that does not undergo electrolysis; it won't conduct electricity and won't be decomposed by it, either in solution or when molten. Sugar, for example, is a non-electrolyte.
Exception: covalent compounds that are electrolytes in solution (because they react with water to form ions) — these include acids and ammonia solution.
💡 These experiments give an easy test of whether a substance is ionic: if it undergoes electrolysis (molten or in solution), it must contain ions; if it doesn't undergo electrolysis, it doesn't contain ions.

🔬 Practical (Activity 7): Quantitative electrolysis — how current affects the amount of product

Quantitative means related to numbers: investigate how the amount of hydrogen produced at the cathode in the electrolysis of sodium chloride solution changes with the current.
Apparatus (Figure 10.11): the electrolytic cell as in Figure 10.6, but with a variable resistor and an ammeter in the circuit (to vary and measure the current), plus a gas burette or measuring cylinder to measure the volume of gas produced.

Method: 1. Set up the apparatus; pour 50 cm³ of concentrated sodium chloride solution into the glass tube. 2. Place a gas burette (or measuring cylinder) filled with sodium chloride solution over the cathode. 3. Turn on the power/connect the battery and set the current to 0.2 A using the variable resistor. 4. Take the initial reading on the gas burette after turning the power on (the current must be flowing so you can see what the current is). 5. Start the timer; after 5 minutes stop the timer and note the final burette reading. 6. Repeat for currents of 0.4 A, 0.6 A, 0.8 A and 1.0 A; repeat each experiment to get more reliable results.

Example results:

Current / A	Volume of hydrogen produced / cm³
0.20	7.0
0.40	13.9
0.60	15.1 (anomalous)
0.80	28.0
1.00	34.9

Graph (Figure 10.12): plot volume of hydrogen (y) against current (x). Current is continuous data, so draw a line of best fit; the reading at 0.6 A is an anomalous point (too low — the current used may have been too low or some gas escaped) and is excluded from the line of best fit.
The line of best fit is a straight line through the origin, showing a proportional relationship: the volume of gas produced is directly proportional to the current passed through the solution.

✅ Quick check

Q: Why does the bulb in the molten lead(II) bromide circuit light up only after the solid has melted? A: Solid PbBr₂ is a giant ionic lattice: the Pb²⁺ and Br⁻ ions are locked in place and cannot move, so no charge flows. When molten, the ions are free to move and carry the current (bulb lights), and electrolysis occurs.
Q: Write the half-equations for the electrolysis of molten sodium chloride and state which is oxidation and which reduction. A: Cathode: Na⁺ + e⁻ → Na (reduction — gain of electrons). Anode: 2Cl⁻ → Cl₂ + 2e⁻ (oxidation — loss of electrons).
Q: Why is hydrogen produced at the cathode when concentrated sodium chloride solution is electrolysed, rather than sodium? A: Both Na⁺ and H⁺ (from water) are attracted to the cathode, but sodium is very reactive — it is very difficult to add an electron to Na⁺. Hydrogen is less reactive, so H⁺ is discharged instead: 2H⁺(aq) + 2e⁻ → H₂(g).
Q: Copper(II) sulfate solution is electrolysed with inert electrodes. What forms at each electrode and what happens to the solution over time? A: Copper at the cathode (Cu²⁺(aq) + 2e⁻ → Cu(s)) and oxygen at the anode (4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻). The blue colour fades to colourless as Cu²⁺ is used up, and the solution becomes acidic / turns into dilute sulfuric acid (H⁺ ions build up).
Q: In the electrolysis of dilute sulfuric acid, why is the volume of hydrogen twice the volume of oxygen, and why is it actually rather more than twice? A: 4 electrons produce 1 molecule of O₂ but 2 molecules of H₂ (twice the number of molecules occupy twice the volume). In practice more hydrogen is collected because oxygen is more soluble in water than hydrogen — saturating the water with oxygen first gives better results.

UNIT 2 · Inorganic Chemistry

Ch 11 The Alkali Metals

Chapter 11: The Alkali Metals

Group 1 of the Periodic Table — lithium, sodium, potassium, rubidium, caesium and francium — forms a family of extremely reactive, soft metals that are all stored away from air and water. They all have one outer electron, react similarly with water (giving an alkaline solution of the metal hydroxide plus hydrogen), and get more reactive as you go down the group. (The section "Explaining the increase in reactivity" is CHEMISTRY ONLY; the rest is Double Award content.)

Learning objectives: understand how the similarities in the reactions of these elements with water provide evidence for their recognition as a family of elements • understand how the differences between their reactions with air and water provide evidence for the trend in reactivity in Group 1 (CHEMISTRY ONLY) • use knowledge of trends in Group 1 to predict the properties of other alkali metals • explain the trend in reactivity in Group 1 in terms of electronic configurations (CHEMISTRY ONLY).

Introduction: the Group 1 elements

The elements in Group 1 of the Periodic Table are called the alkali metals: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr).
Francium is radioactive: one of its isotopes is produced during the radioactive decay of uranium-235, but it is extremely short-lived. At any one time there is only about 20–30 g of francium in the whole of the Earth's crust, and no one has ever seen a piece of francium — but knowing the rest of Group 1 lets you predict what it would be like.
All the alkali metals are very reactive metals that react vigorously with water (Figure 11.2: potassium reacting with water).

Physical properties

Element	Melting point / °C	Boiling point / °C	Density / g cm⁻³
Li	181	1342	0.53
Na	98	883	0.97
K	63	760	0.86
Rb	39	686	1.53
Cs	29	669	1.88

Melting and boiling points are very low for metals, and they decrease as you move down the group.
Extension work: the melting points decrease because the atoms get bigger down the group — in the metallic lattice the positive ions are further from the delocalised electrons (e.g. in caesium compared with sodium), so there is weaker electrostatic attraction (weaker metallic bonding).
Densities tend to increase down the group, although not regularly.
Lithium, sodium and potassium are all less dense than water (0.53, 0.97 and 0.86 g cm⁻³ < 1.0 g cm⁻³), so they float on water.
The metals are very soft — easily cut with a knife — and become softer as you go down the group.
They are shiny and silver when freshly cut, but tarnish very quickly on exposure to air (they react with oxygen in the air to form oxides).

Storage and handling

All the alkali metals are extremely reactive, and they get more reactive as you go down the group. They all react quickly with oxygen in the air (forming oxides) and react rapidly with water to form strongly alkaline solutions of the metal hydroxides — this is why they are called the alkali metals.
Lithium, sodium and potassium are stored under oil to stop them reacting with the oxygen or water vapour in the air.
Rubidium and caesium are so reactive that they have to be stored in sealed glass tubes to stop any possibility of oxygen getting at them.
⚠️ Great care must be taken not to touch any of these metals with bare fingers — there could be enough sweat on your skin to give a reaction, producing lots of heat and a very corrosive metal hydroxide.

💡 Exam tip: when asked to describe observations, the book's hint says it is better to write "fizzing / bubbling / effervescing" rather than "a gas is given off", because the fizzing/bubbling/effervescing is what you actually see.

A family of elements

There are two reasons these elements are put in Group 1:

They all have one electron in their outer shell. Electronic configurations:

lithium   2, 1
sodium    2, 8, 1
potassium 2, 8, 8, 1

They have similar chemical properties, for example: - they all react with water (in the same way) to form a hydroxide with the formula MOH (e.g. LiOH, NaOH) and hydrogen; - they react with oxygen to form an oxide with the formula M₂O (e.g. Na₂O, K₂O); - they react with halogens to form compounds with the formula MX (e.g. LiCl, KBr); - they form ionic compounds containing an M⁺ ion (e.g. Na⁺, K⁺). - (In these formulae, 'M' simply represents any one of the alkali metals and 'X' any one of the halogens.) - The Group 1 elements react in very similar ways because they all have the same number of electrons in the outer shell (one) — chemical properties depend on the number of outer-shell electrons, so reason 2 is really just a consequence of reason 1. - 💡 DID YOU KNOW? Historically, the alkali metals were put in the same group because of similarities in properties — before anyone knew anything about electrons!

Reactions with water

All these metals react in the same way with water to produce a metal hydroxide and hydrogen:

alkali metal + water → alkali metal hydroxide + hydrogen
2M + 2H₂O → 2MOH + H₂

(MOH makes sense because the charge on the hydroxide ion is 1−.) - The main difference between the reactions is how quickly they happen: as you go down the group, the metals become more reactive and the reactions occur more rapidly.

With sodium (typical reaction)

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Main observations (Figure 11.5): - The sodium floats — because it is less dense than water. - The sodium melts into a ball — its melting point is low and a lot of heat is produced by the reaction. - There is fizzing — because hydrogen gas is produced. - The sodium moves rapidly around the surface of the water — the hydrogen is not given off symmetrically around the ball, so the sodium is pushed around the surface like a hovercraft. - A white trail is formed which dissolves in the water — this is the sodium hydroxide. - The piece of sodium gets smaller and eventually disappears — it is used up in the reaction. - Test the solution with universal indicator: it goes blue (alkaline) — the metal hydroxide is alkaline (the solution contains OH⁻ ions).

💡 Exam tips (from the book's Hints): - Strictly speaking, most of the time the sodium is reacting it is present as molten sodium, not solid sodium — but writing (l) for the state symbol of sodium has the potential to confuse an examiner and is probably best avoided (write Na(s)). - When asked to compare reactions, use phrases such as: "fizzes more/less vigorously", "moves around more quickly/slowly", "disappears more quickly/takes longer to disappear" and say how you can see it (e.g. the lithium fizzes more slowly, moves around more slowly, or takes longer to disappear).

With lithium

2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g)

The reaction is very similar to sodium's, except that it is slower.
Lithium doesn't melt — its melting point is higher and the heat isn't produced so quickly, so the heat isn't enough to melt it.

With potassium

2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

Potassium's reaction is faster than sodium's: enough heat is produced to ignite the hydrogen, which burns with a lilac flame. The reaction often ends with the potassium spitting around and exploding.
The key difference from sodium: with potassium the hydrogen bursts into flames; with sodium it usually does not.

With rubidium and caesium

These react even more violently than potassium, and the reaction can be explosive. Rubidium hydroxide and caesium hydroxide are formed.
⚠️ Safety Note (CHEMISTRY ONLY): the reactions of rubidium and caesium would be too hazardous to attempt in school — they would explode.

Explaining the increase in reactivity — CHEMISTRY ONLY

As you go down the group the metals become more reactive.
In all these reactions the metal atoms are losing electrons and forming metal ions in solution, e.g.:

Na(s) → Na⁺(aq) + e⁻

The differences between the reactions depend on how easily the outer electron is lost in each case — which depends on how strongly it is attracted to the nucleus in the original atom. (The nucleus is positive because it contains protons, so it attracts the negative electrons.)
As we move down the group, the atoms have more shells of electrons and get bigger — a sodium atom is bigger than a lithium atom, a potassium atom bigger than a sodium atom. The outer electron (the one lost in the reaction) is therefore further from the nucleus.
Because it is further from the nucleus it is less strongly attracted by the nucleus and therefore more easily lost — so reactivity increases down the group.
The electrons released by the metal are gained by the water molecules, producing hydroxide ions and hydrogen gas:

2H₂O(l) + 2e⁻ → 2OH⁻(aq) + H₂(g)

📊 Key diagram (Figure 11.7): lithium vs sodium atoms

   Li:  nucleus (3+)         Na:  nucleus (11+)
         ●  (inner shell)           ●   (inner shells)
           ○ ← outer electron          ○   ○ ← outer electron
   outer electron close to nucleus →  outer electron further from
   strongly attracted → less easily    nucleus → more easily lost
   lost                                 (so Na is more reactive)

Reactions of the alkali metals with the air

Lithium, sodium and potassium are stored in oil because they react with the air. A piece of sodium taken out of the oil usually has a crust on the outside and isn't shiny unless freshly cut.
When a piece of sodium is cut, the fresh surface is shiny but it tarnishes rapidly as the freshly exposed sodium reacts with oxygen in the air.
With lithium the same thing happens but it tarnishes more slowly (lithium reacts more slowly than sodium). A freshly cut piece of potassium tarnishes extremely rapidly — more quickly than sodium. In this way we can see again that potassium is more reactive than sodium, which is more reactive than lithium.
In each case the metal reacts with oxygen in the air to form an oxide with the formula M₂O.
If we heat each metal in the air with a Bunsen burner, we get a much more vigorous reaction (it becomes harder to tell which is most reactive because all the reactions are so rapid):
Lithium burns with a red flame to form lithium oxide.
Sodium burns with a yellow flame to form sodium oxide.
Potassium burns with a lilac flame to form potassium oxide.
The equation for all these reactions is:

4M(s) + O₂(g) → 2M₂O(s)

In each case the product is a white powder — the alkali metal oxide.
Extension work: when the alkali metals react with air, other types of oxide can also form — peroxides and superoxides (all still contain the M⁺ ion; it is the oxygen part which is different). (Book's formula subscripts for these were garbled in the OCR; know the names and that M⁺ is present.) Lithium can also form a nitride, Li₃N.

Compounds of the alkali metals

All Group 1 metal ions are colourless, so their compounds are colourless or white unless combined with a coloured negative ion.
Examples of coloured compounds: potassium dichromate(VI) is orange because the dichromate(VI) ion is orange; potassium manganate(VII) is purple because the manganate(VII) ion is purple.
Group 1 compounds are typical ionic solids and are mostly soluble in water.

Summarising the main features of the Group 1 elements

Group 1 elements: - are metals; - are soft, with melting points and densities which are very low for metals; - have to be stored out of contact with air or water; - react rapidly with air to form coatings of the metal oxide; - react with water to produce an alkaline solution of the metal hydroxide and hydrogen gas; - show an increase in reactivity as you go down the group; - form compounds in which the metal has a 1+ ion (M⁺); - have mainly white/colourless compounds which dissolve to produce colourless solutions.

Predicting the properties of francium

As we move down a group the properties of the elements change gradually, so knowing the properties of most elements in a group lets us predict the properties of elements we don't know (you can't actually observe any of these for francium — it's far too rare and short-lived).
We can predict that francium:
is very soft;
has a melting point around room temperature;
has a density probably just over 2 g/cm³;
is a silvery metal, but tarnishes almost instantly in air;
reacts violently with water to give francium hydroxide and hydrogen;
is more reactive than caesium;
has a hydroxide, FrOH, which is soluble in water and forms a strongly alkaline solution;
forms compounds that are white/colourless, dissolving in water to give colourless solutions.
Graphical method (Figure 11.9): plot melting point against atomic number for the alkali metals and draw a line of best fit; extending the line to atomic number 87 (francium) predicts a melting point of about 22 °C (other predictions give values between 21 °C and 27 °C).

✅ Quick check

Q: Explain why lithium, sodium and potassium are placed in the same group of the Periodic Table. A: They all have one electron in their outer shell (configurations 2,1 / 2,8,1 / 2,8,8,1), and they have similar chemical properties — reacting with water to give MOH + H₂, forming M₂O oxides, MX halides and M⁺ ions. The similar properties are a consequence of the same number of outer-shell electrons.
Q: Describe what you would observe when a piece of sodium is dropped onto water, and write the equation. A: It floats; melts into a ball; fizzes (hydrogen gas evolved); moves rapidly around the surface like a hovercraft; a white trail (sodium hydroxide) dissolves in the water; the piece gets smaller and disappears; universal indicator turns blue (alkaline solution). Equation: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g).
Q: Arrange Li, Na and K in order of increasing reactivity with water, and explain the trend in terms of electronic configurations. (CHEMISTRY ONLY) A: Li < Na < K (potassium most reactive; may ignite the hydrogen with a lilac flame). Going down the group the atoms have more electron shells and get bigger, so the outer electron is further from the nucleus, less strongly attracted, and more easily lost — so reactivity increases.
Q: What flame colours do lithium, sodium and potassium give when burned in air, and what is the general equation? A: Lithium — red flame; sodium — yellow flame; potassium — lilac flame. 4M(s) + O₂(g) → 2M₂O(s) (white powder metal oxide).
Q: Predict four properties of francium using the Group 1 trends. A: Very soft; melting point around room temperature (≈22 °C by extrapolation); density just over 2 g/cm³; silvery but tarnishes almost instantly in air; reacts violently with water (more reactive than caesium) giving FrOH (soluble, strongly alkaline) and hydrogen; white/colourless compounds giving colourless solutions.

Ch 12 The Halogens

Chapter 12: The Halogens (Group 7)

The halogens are the reactive non-metals in Group 7: fluorine, chlorine, bromine and iodine. They have 7 outer electrons, exist as diatomic molecules, and become less reactive down the group. They form salts with metals and can displace less reactive halogens from their compounds.

THE HALOGEN ELEMENTS

Group 7 elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I) (+ astatine).
They have 7 outer-shell electrons → they need 1 electron to gain a full outer shell → they form −1 ions (e.g. Cl⁻) or share electrons covalently.
They exist as diatomic molecules: F₂, Cl₂, Br₂, I₂.

PHYSICAL PROPERTIES (trends down the group)

Element	State at room temp	Colour
Fluorine	Gas	Pale yellow
Chlorine	Gas	Green-yellow
Bromine	Liquid	Red-brown
Iodine	Solid	Grey-black (purple vapour)
- Down the group: melting/boiling points increase (F gas → I solid); colour gets darker.

REACTIVITY

Reactivity DECREASES down the group (as atoms get bigger, the outer electrons are further from the nucleus and less able to attract an extra electron).
KEY POINT: fluorine is the most reactive non-metal; iodine is the least reactive of these.
All halogens are toxic and disinfectants.

REACTIONS

With metals: form salts (halides) — e.g. sodium + chlorine → sodium chloride (2Na + Cl₂ → 2NaCl).
With hydrogen: form hydrogen halides (e.g. HCl) — which dissolve in water to form acids.
With water (chlorine): chlorine + water → hydrochloric acid + chloric(I) acid (HClO) — this is why chlorine is added to water supplies (kills bacteria).

DISPLACEMENT REACTIONS

A more reactive halogen displaces a less reactive one from its salt solution.
Chlorine displaces bromine and iodine from their salts.
Bromine displaces iodine but NOT chlorine.
Iodine displaces neither.
Examples:
Cl₂ + 2KBr → 2KCl + Br₂ (solution turns orange/brown).
Cl₂ + 2KI → 2KCl + I₂ (solution turns brown/black).
Br₂ + 2KI → 2KBr + I₂.
Observations: the solution changes colour (orange for bromine, brown for iodine).

USES OF HALOGENS

Chlorine: sterilising drinking water and swimming pools; making bleach, PVC (as HCl).
Iodine: antiseptic (iodine solution), essential nutrient (thyroid).
Fluorine: in toothpaste (fluoride prevents tooth decay).
Bromine: flame retardants, photography.

✅ Quick check

How many outer electrons do halogens have? (7.)
What is the trend in reactivity down Group 7? (Reactivity decreases.)
What happens when chlorine is bubbled into potassium bromide solution? (Chlorine displaces bromine — the solution turns orange/brown: Cl₂ + 2KBr → 2KCl + Br₂.)
Give one use of chlorine. (Sterilising water, making bleach.)
What are the states of chlorine, bromine and iodine at room temperature? (Gas, liquid, solid.)

Ch 13 Gases in the Atmosphere

Chapter 13: Gases in the Atmosphere

The atmosphere is mainly nitrogen and oxygen, with small amounts of other gases. This chapter covers the composition of air, the properties and uses of the main gases, the greenhouse effect and climate change, and the tests for gases (oxygen, carbon dioxide, hydrogen, chlorine).

COMPOSITION OF THE ATMOSPHERE (memorise the percentages)

Nitrogen (N₂): ~78%
Oxygen (O₂): ~21%
Argon (and other noble gases): ~0.9%
Carbon dioxide (CO₂): ~0.04%
Water vapour (variable).
(The atmosphere also contains trace gases.)

NITROGEN

Unreactive (strong triple bond in N₂) — used as a blanket/inert atmosphere (e.g. in food packaging to stop oxidation; in light bulbs).
Used to make ammonia (Haber process) and fertilisers.

OXYGEN

Supports combustion (burning) and respiration.
Used in hospitals (oxygen therapy), welding (with acetylene), steel making.

CARBON DIOXIDE

Used in fizzy drinks, fire extinguishers, photosynthesis.
Testing for CO₂: bubble the gas through limewater (calcium hydroxide solution) → turns milky/cloudy (white precipitate).

NOBLE GASES (Group 0)

Very unreactive (full outer shells) — used in discharge tubes (neon signs glow), argon in welding, helium in balloons (less dense than air).

THE GREENHOUSE EFFECT AND CLIMATE CHANGE

Greenhouse gases: carbon dioxide, methane (and water vapour). They trap heat in the atmosphere (absorb infrared radiation) — the natural greenhouse effect keeps the Earth warm enough for life.
Enhanced greenhouse effect: human activity (burning fossil fuels, deforestation, farming) increases CO₂ and methane → more heat trapped → global warming → climate change (rising sea levels, extreme weather, habitat loss).
Reducing it: renewable energy, energy efficiency, reforestation, reducing meat/dairy, international agreements (e.g. Paris Agreement).

POLLUTANT GASES FROM BURNING FUELS

Carbon monoxide (CO): incomplete combustion — toxic (binds to haemoglobin).
Sulfur dioxide (SO₂): from sulfur impurities in fuels → acid rain.
Nitrogen oxides (NOx): from high-temperature combustion in engines → acid rain and smog.
Particulates (soot): from diesel engines — damage lungs.
Catalytic converters and clean fuels reduce these.

TESTS FOR GASES (memorise — very common exam questions)

Gas	Test	Positive result
Oxygen	Glowing splint	Relights the splint
Hydrogen	Burning splint	Squeaky pop (burns with a pop)
Carbon dioxide	Bubble through limewater	Turns milky/cloudy
Chlorine	Damp blue litmus paper	Bleaches it (turns white)
Ammonia	Damp red litmus paper	Turns blue (alkaline)

SHOWING THAT AIR CONTAINS ABOUT ONE-FIFTH OXYGEN (practicals)

Air is about 1/5 oxygen (21%). Three classic practicals prove it: 1. 🔬 Using phosphorus: float a dish of white phosphorus in a bell jar of air, sealed over water. The phosphorus burns, using up the oxygen; the water level rises to about 1/5 of the jar. 2. 🔬 Using the rusting of iron: moist iron filings in a sealed tube over water rust, removing oxygen; the water rises to about 1/5 of the tube. 3. 🔬 Using copper: pass air over heated copper turnings — the copper reacts with oxygen (turns black, copper oxide); the volume of air falls by ~1/5. (Repeat until no more reaction.)
Conclusion: ~20% of air is oxygen (the rest is mostly nitrogen, which does not react).

THE COMBUSTION OF ELEMENTS IN OXYGEN

Elements burn in oxygen to form oxides:
Carbon burns brightly → carbon dioxide (CO₂ — turns limewater milky).
Sulfur burns with a blue flame → sulfur dioxide (SO₂ — pungent, acidic).
Magnesium burns with a dazzling white light → magnesium oxide (white powder: 2Mg + O₂ → 2MgO).

✅ Quick check

Give the % composition of air. (78% nitrogen, 21% oxygen, ~0.9% argon, ~0.04% CO₂.)
How do you test for oxygen, hydrogen and carbon dioxide? (O₂: relights a glowing splint; H₂: squeaky pop; CO₂: turns limewater milky.)
Name two greenhouse gases. (Carbon dioxide, methane.)
What causes acid rain? (Sulfur dioxide and nitrogen oxides from burning fuels.)
Why is nitrogen used in food packaging? (It is unreactive — prevents oxidation/spoilage.)

Ch 14 Reactivity Series

Chapter 14: The Reactivity Series

The reactivity series orders metals by how readily they react. More reactive metals displace less reactive ones, react faster with water/acid, and need more energy to extract from their ores. This chapter covers the series itself, reactions with water and acids, and displacement.

THE REACTIVITY SERIES (memorise the order)

POTASSIUM     K     (most reactive)
SODIUM        Na
LITHIUM       Li
CALCIUM       Ca
MAGNESIUM     Mg
ALUMINIUM     Al
(ZINC)        Zn
IRON          Fe
(LEAD)        Pb
COPPER        Cu
SILVER        Ag
GOLD          Au     (least reactive)

KEY POINT: the more reactive the metal, the more easily it loses electrons to form positive ions.

REACTIONS WITH WATER

Very reactive metals (K, Na, Li, Ca): react vigorously with cold water, giving the metal hydroxide + hydrogen.
Sodium + water → sodium hydroxide + hydrogen: 2Na + 2H₂O → 2NaOH + H₂ (sodium fizzes, melts into a ball, moves on the surface).
Potassium is even more violent (lilac flame); lithium is gentler; calcium fizzes and goes cloudy (insoluble hydroxide).
Less reactive metals (Mg, Zn, Fe): react very slowly with cold water; magnesium reacts with steam (heat) to give magnesium oxide + hydrogen: Mg + H₂O → MgO + H₂.
Copper and below: no reaction with water or steam.

REACTIONS WITH ACIDS

Metals ABOVE hydrogen in the series react with dilute acids to give a salt + hydrogen:
metal + acid → salt + hydrogen
Example: magnesium + hydrochloric acid → magnesium chloride + hydrogen: Mg + 2HCl → MgCl₂ + H₂ (bubbles of hydrogen — squeaky pop).
Rate of bubbling shows reactivity: Mg reacts fast, Zn slower, Fe slowest; Cu does not react with dilute acids.
KEY POINT: the speed of reaction with acid ranks metals in the reactivity series.

DISPLACEMENT REACTIONS

A more reactive metal displaces a less reactive metal from a solution of its salt.
Example: iron + copper sulfate → iron sulfate + copper: Fe + CuSO₄ → FeSO₄ + Cu (blue solution fades, brown copper deposits).
Zinc displaces copper and iron; copper cannot displace zinc or iron.
This is used in thermit welding (aluminium displaces iron from iron oxide) and explains rusting protection (sacrificial protection with zinc).

EXTRACTION OF METALS (see Ch 15 for detail)

The more reactive the metal, the harder it is to extract (stronger bonds in the ore):
Metals below carbon (Zn, Fe, Cu...) — extract by reduction with carbon (heating the oxide with carbon/coke).
Metals above carbon (K, Na, Ca, Mg, Al) — extract by electrolysis of the molten compound.
Gold/silver — found native (uncombined) in the ground.
💡 Exam tip: carbon can only reduce metals BELOW it in the series; the series is used to choose the extraction method.

PREVENTING RUSTING (and displacement with copper)

Rusting — the corrosion of iron/steel needing BOTH oxygen AND water (test: nail in dry air doesn't rust; nail in boiled (air-free) water doesn't rust; nail in air + water rusts).
Methods of prevention:
Painting, oiling, greasing, plastic coating — barrier stops air/water reaching the iron.
Galvanising — coating with zinc: even if scratched, zinc corrodes instead (zinc is more reactive — sacrificial protection).
Sacrificial protection with a more reactive metal (zinc blocks attached to ship hulls/pipelines).
Alloying — making stainless steel (with chromium and nickel) prevents rusting.
Displacement example — copper and silver nitrate: copper + silver nitrate → copper nitrate + silver: Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag (silver crystals deposit, solution turns blue). Copper is more reactive than silver, so it displaces it.

✅ Quick check

Write the reactivity series from potassium to gold. (K, Na, Li, Ca, Mg, Al, Zn, Fe, Pb, Cu, Ag, Au.)
What happens when sodium reacts with water? (Fizzes, melts, moves — sodium hydroxide + hydrogen.)
Write the equation for magnesium + hydrochloric acid. (Mg + 2HCl → MgCl₂ + H₂.)
What is a displacement reaction? (A more reactive metal displaces a less reactive one from its salt solution.)
Why must aluminium be extracted by electrolysis? (It is above carbon — carbon cannot reduce its oxide.)

Ch 15 Extraction and Uses of Metals

Chapter 15: Extraction and Uses of Metals

Metals are extracted from their ores. The method depends on the metal's position in the reactivity series: unreactive metals occur native; metals below carbon are extracted by reduction with carbon; metals above carbon need electrolysis. This chapter covers the extraction of iron (blast furnace), aluminium (electrolysis), and the uses of metals.

CHOOSING THE EXTRACTION METHOD (KEY IDEA)

Metal	Position	Method
Gold, silver	Bottom of series	Native (found uncombined) — mined directly
Zinc, iron, lead, copper	Below carbon	Reduction with carbon/coke (heating the oxide with carbon)
Potassium → aluminium	Above carbon	Electrolysis of the molten compound (carbon can't displace them)

EXTRACTION OF IRON — THE BLAST FURNACE

Raw materials: iron ore (haematite, Fe₂O₃), coke (carbon), limestone (calcium carbonate) and hot air.
Process: 1. Hot air is blown in the bottom — coke burns: C + O₂ → CO₂ (exothermic — huge heat). 2. CO₂ reacts with more coke to form carbon monoxide: CO₂ + C → 2CO. 3. CO reduces the iron oxide: Fe₂O₃ + 3CO → 2Fe + 3CO₂. (Fe³⁺ is reduced to Fe.) 4. Limestone removes impurities (sand) as slag (calcium silicate): CaCO₃ → CaO + CO₂, then CaO + SiO₂ → CaSiO₃. 5. Molten iron (tapped from the bottom) and molten slag (used in road building) are removed.
KEY POINT: the blast furnace produces iron — the carbon reduces the iron oxide (removes oxygen).
Steel: iron is converted to steel by removing impurities and adding controlled amounts of carbon (and other metals) — steel is harder/stronger than pure iron.
Different kinds of steel (different uses):
Mild steel (low carbon) — easily shaped: car bodies, wires, girders.
High-carbon steel — hard but brittle: tools, blades, drill bits.
Stainless steel (with chromium + nickel) — rust-resistant: cutlery, sinks, surgical instruments.

EXTRACTION OF ALUMINIUM — ELECTROLYSIS

Aluminium is extracted by electrolysis of molten aluminium oxide (bauxite):
Bauxite is purified to aluminium oxide, then dissolved in molten cryolite (lowers the melting point — saves energy).
Cathode (−): aluminium ions gain electrons → aluminium metal: Al³⁺ + 3e⁻ → Al.
Anode (+): oxide ions lose electrons → oxygen gas: 2O²⁻ → O₂ + 4e⁻.
The oxygen reacts with the carbon anodes, burning them away (they must be replaced).
Why electrolysis? Aluminium is above carbon — carbon cannot reduce its oxide. Electrolysis is expensive (huge electricity) → aluminium is recycled to save ~95% of the energy.

EXTRACTION OF COPPER

Copper is extracted from copper sulfide ores by smelting (heating in air), then purified by electrolysis (impure copper anode, pure copper cathode — the copper ions move to the cathode).
Recycling copper saves energy and reduces mining.

USES OF METALS (linked to their properties)

Aluminium: light, strong, corrosion-resistant → aircraft, drink cans, foil, window frames.
Copper: excellent conductor of electricity and heat, malleable → electrical wiring, pipes, cooking pans.
Iron/steel: strong, cheap → buildings, cars, bridges, tools.
Gold: unreactive, shiny, good conductor → jewellery, electronics contacts, currency.
Zinc: used to galvanise (coat) iron/steel to prevent rusting (sacrificial protection).

✅ Quick check

What determines the extraction method? (Position in the reactivity series.)
What reduces iron oxide in the blast furnace? (Carbon monoxide: Fe₂O₃ + 3CO → 2Fe + 3CO₂.)
Why is limestone added to the blast furnace? (To remove sand impurities as slag.)
Why is aluminium extracted by electrolysis? (It is above carbon in the reactivity series — carbon can't reduce its oxide.)
What happens at the cathode during aluminium extraction? (Al³⁺ + 3e⁻ → Al — aluminium metal forms.)
Why is copper used in wiring? (Excellent conductor of electricity.)

Ch 16 Acids, Alkalis and Titrations

Chapter 16: Acids, Alkalis and Titrations

Acids produce H⁺ ions in water (pH < 7); alkalis produce OH⁻ ions (pH > 7). This chapter covers the pH scale, indicators, neutralisation, strong vs weak acids, and the titration method for finding concentrations.

ACIDS AND ALKALIS

Acid — a substance that produces hydrogen ions (H⁺) when dissolved in water (pH below 7).
Examples: hydrochloric acid (HCl), sulfuric acid (H₂SO₄), nitric acid (HNO₃), citric acid, ethanoic acid.
Alkali — a soluble base that produces hydroxide ions (OH⁻) in water (pH above 7).
Examples: sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide, ammonia solution.
Base — a substance that neutralises an acid (metal oxides and hydroxides); an alkali is a base that dissolves in water.
Neutral — pH 7 (pure water).

THE pH SCALE AND INDICATORS

pH scale: 0–6 acid, 7 neutral, 8–14 alkali. Lower pH = stronger acid; higher pH = stronger alkali.
Universal indicator: gives a colour for every pH (red = strong acid, green = neutral, purple = strong alkali).
Litmus: red in acid, blue in alkali.
pH can be measured with a pH meter (more accurate) or pH paper.
The lower the pH, the higher the concentration of H⁺ ions.

NEUTRALISATION

Neutralisation — the reaction of an acid with a base to form a salt + water: acid + base → salt + water H⁺(aq) + OH⁻(aq) → H₂O(l)
Examples: HCl + NaOH → NaCl + H₂O; H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O.
KEY POINT: the ionic equation for ALL neutralisations is H⁺ + OH⁻ → H₂O.

STRONG vs WEAK ACIDS

Strong acid — fully ionises in water (every acid molecule splits into H⁺ and its anion): HCl → H⁺ + Cl⁻. Low pH (e.g. 1).
Weak acid — partially ionises (only some molecules split): ethanoic acid CH₃COOH ⇌ CH₃COO⁻ + H⁺. Higher pH (e.g. 4–5).
KEY POINT: pH depends on how much the acid ionises, not just how concentrated it is. A concentrated weak acid can have a higher pH than a dilute strong acid.

TITRATIONS (finding concentrations)

Titration — a method to find the exact volume of one solution needed to react completely with a known volume of another.
Method (acid–alkali titration): 1. Use a pipette to measure an exact volume of alkali (e.g. 25 cm³) into a conical flask. 2. Add a few drops of indicator (e.g. phenolphthalein or methyl orange). 3. Fill a burette with the acid. 4. Add the acid slowly, swirling, until the indicator just changes colour (the end point — neutralisation). 5. Record the titre (volume of acid used); repeat for accuracy and take the mean.
Calculation: moles = concentration (mol/dm³) × volume (dm³) n = c × V (remember: 1 dm³ = 1000 cm³) Then use the balanced equation's mole ratio to find the unknown concentration.
Example: 25.0 cm³ of NaOH (0.10 mol/dm³) needs 20.0 cm³ of HCl. Moles NaOH = 0.10 × 0.025 = 0.0025 mol. Ratio 1:1 → moles HCl = 0.0025 → concentration HCl = 0.0025 ÷ 0.020 = 0.125 mol/dm³.
Washing: the flask is rinsed with distilled water (not the solution — would change the moles).

✅ Quick check

What ions do acids and alkalis produce in water? (H⁺ and OH⁻.)
What is neutralisation? (Acid + base → salt + water; H⁺ + OH⁻ → H₂O.)
Distinguish strong and weak acids. (Strong = fully ionised; weak = partially ionised.)
What apparatus is used in a titration? (Pipette, burette, conical flask, indicator.)
Calculate the moles in 50 cm³ of 0.2 mol/dm³ solution. (0.2 × 0.05 = 0.01 mol.)

Ch 17 Acids, Bases and Salt Preparations

Chapter 17: Acids, Bases and Salt Preparations

This chapter covers the reactions of acids (with metals, bases, carbonates) and the three methods for making salts: acid + excess insoluble solid, acid + alkali (titration), and precipitation. You must know which method suits which salt.

SALTS

Salt — the compound formed when the H⁺ of an acid is replaced by a metal ion (or ammonium ion).
The salt's name comes from the acid: hydrochloric acid → chloride; sulfuric acid → sulfate; nitric acid → nitrate.
Examples: HCl + NaOH → NaCl; H₂SO₄ + CuO → CuSO₄ + H₂O.

THE FOUR REACTIONS OF ACIDS

Acid + metal → salt + hydrogen - Mg + 2HCl → MgCl₂ + H₂ (only metals above hydrogen in the reactivity series).
Acid + base (metal oxide/hydroxide) → salt + water - CuO + H₂SO₄ → CuSO₄ + H₂O (neutralisation).
Acid + alkali → salt + water - HCl + NaOH → NaCl + H₂O.
Acid + carbonate → salt + water + carbon dioxide - CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ (fizzes — test with limewater). - Na₂CO₃ + H₂SO₄ → Na₂SO₄ + H₂O + CO₂.

MAKING SALTS — THE THREE METHODS

Method 1: ACID + EXCESS INSOLUBLE SOLID (for salts of metals that don't react with water — e.g. copper sulfate from copper oxide)

Warm the acid (speeds the reaction).
Add the insoluble base (e.g. copper oxide) in excess — keep adding until no more dissolves (some solid remains).
Filter to remove the excess solid.
Evaporate/crystallise the filtrate to get the salt crystals. - Example: CuO + H₂SO₄ → CuSO₄ + H₂O → blue copper sulfate crystals.

Method 2: ACID + ALKALI (TITRATION) (for soluble salts where both reactants dissolve — e.g. sodium chloride)

Use a titration (Chapter 16) to find the exact volumes that neutralise each other (using an indicator).
Repeat WITHOUT the indicator, mixing the exact volumes.
Evaporate/crystallise to get the pure salt. - Why not excess? Both reactants are soluble — you can't filter out excess; you need exact amounts.

Method 3: PRECIPITATION (for INSOLUBLE salts — e.g. lead iodide, barium sulfate)

Mix two soluble solutions whose ions form an insoluble salt (the precipitate).
Example: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) (yellow precipitate) + 2KNO₃(aq).
Method: mix the two solutions → the precipitate forms → filter → wash with distilled water → dry.
Example use: making barium sulfate (used in X-rays).

SOLUBILITY RULES (quick guide)

Most soluble: all nitrates; all sodium, potassium, ammonium salts; most chlorides (except AgCl, PbCl₂); most sulfates (except BaSO₄, PbSO₄).
Insoluble: most carbonates and hydroxides (except Na, K, ammonium).

✅ Quick check

Name the four reactions of acids. (Acid + metal/base/alkali/carbonate.)
Write the equation for acid + carbonate. (Acid + carbonate → salt + water + carbon dioxide.)
Which method makes copper sulfate from copper oxide? (Acid + excess insoluble solid — filter, then evaporate/crystallise.)
Why can't you use the "excess solid" method for sodium chloride? (Both reactants are soluble — nothing to filter; use titration.)
How do you make an insoluble salt? (Precipitation — mix two solutions, filter, wash, dry.)

Ch 18 Chemical Tests

Chapter 18: Chemical Tests

This chapter collects all the tests you must know: flame tests for metal ions, tests for cations (with sodium hydroxide and ammonia), tests for anions (carbonates, sulfates, halides, nitrates), and the gas tests. These are guaranteed marks — memorise the reagents AND the positive results.

FLAME TESTS (for metal cations)

Dip a clean nichrome wire loop in the solid/solution and hold in a blue Bunsen flame: | Metal ion | Flame colour | |---|---| | Lithium (Li⁺) | Crimson/red | | Sodium (Na⁺) | Yellow/orange | | Potassium (K⁺) | Lilac | | Calcium (Ca²⁺) | Brick red/orange-red | | Copper (Cu²⁺) | Blue-green/green |
💡 Exam tip: clean the wire in acid between tests.

TESTS FOR CATIONS — WITH SODIUM HYDROXIDE (adding NaOH)

Many metal hydroxides are insoluble → a coloured precipitate forms: | Ion | Result with NaOH | |---|---| | Cu²⁺ | Blue precipitate (Cu(OH)₂) | | Fe²⁺ | Green precipitate (Fe(OH)₂) | | Fe³⁺ | Red-brown precipitate (Fe(OH)₃) | | Zn²⁺ | White precipitate — dissolves in excess NaOH | | Al³⁺ | White precipitate — dissolves in excess NaOH | | Ca²⁺ | White precipitate — insoluble in excess | | NH₄⁺ | No precipitate — warming gives ammonia gas (turns damp red litmus blue) |
💡 Exam tip: "add NaOH → blue precipitate" identifies Cu²⁺; the Zn²⁺/Al³⁺ white precipitates dissolving in excess NaOH distinguish them from Ca²⁺.

TESTS FOR ANIONS

Anion	Test	Positive result
Carbonate (CO₃²⁻)	Add dilute acid	Fizzes — gas turns limewater milky (CO₂)
Sulfate (SO₄²⁻)	Add dilute HCl + barium chloride	White precipitate of BaSO₄
Halides (Cl⁻, Br⁻, I⁻)	Add dilute HNO₃ + silver nitrate	Cl⁻: white precipitate (AgCl); Br⁻: cream (AgBr); I⁻: yellow (AgI)
Nitrate (NO₃⁻)	Add NaOH + aluminium foil, warm	Ammonia gas produced (turns red litmus blue)
- (Add acid first to remove carbonate interference when testing sulfates/halides.)

GAS TESTS (recap — see Ch 13)

Oxygen: glowing splint relights.
Hydrogen: burning splint → squeaky pop.
Carbon dioxide: turns limewater milky.
Chlorine: bleaches damp litmus paper.
Ammonia: turns damp red litmus blue (pungent smell).

TESTING FOR WATER (and the physical test)

Test for pure water (physical): boiling point = 100 °C at standard pressure; freezing point 0 °C.
Chemical test for water: add to anhydrous copper(II) sulfate (white powder) → turns blue if water is present (CuSO₄ + 5H₂O → CuSO₄·5H₂O). Or cobalt chloride paper (blue) → turns pink in water.

✅ Quick check

What are the flame colours for sodium, potassium, lithium and copper? (Yellow, lilac, crimson, blue-green.)
What precipitate forms when NaOH is added to Fe³⁺? (Red-brown.) To Cu²⁺? (Blue.)
How do you test for a sulfate? (Add HCl + barium chloride → white precipitate.)
How do you distinguish Cl⁻, Br⁻ and I⁻? (Silver nitrate: white, cream, yellow precipitates.)
How do you test for ammonium ions? (Add NaOH and warm — ammonia gas turns damp red litmus blue.)

UNIT 3 · Physical Chemistry

Ch 19 Energetics

Chapter 19: Energetics

Chemical reactions either release energy (exothermic) or absorb it (endothermic). This chapter covers the two reaction types, energy level diagrams, bond breaking and making, and how to measure the energy change (calorimetry).

EXOTHERMIC AND ENDOTHERMIC REACTIONS

Exothermic reaction — a reaction that releases heat energy to the surroundings (temperature of the surroundings rises).
Examples: combustion (burning), neutralisation (acid + alkali), respiration, most displacement reactions, adding water to calcium oxide (slaked lime).
Endothermic reaction — a reaction that absorbs heat energy from the surroundings (temperature falls).
Examples: thermal decomposition (e.g. calcium carbonate → calcium oxide + CO₂, heated), photosynthesis, ammonium nitrate dissolving in water, citric acid + sodium hydrogencarbonate (fizzes and gets cold).
💡 Exam tip: "the temperature rose → exothermic; fell → endothermic."

ENERGY LEVEL DIAGRAMS

Exothermic: the products have LESS energy than the reactants (energy is released). The difference is the energy given out (ΔH is negative).
Endothermic: the products have MORE energy than the reactants (energy is absorbed). ΔH is positive.
📊 Diagram: draw two levels — reactants on the left, products on the right; exothermic: products lower, arrow down labelled "energy released"; endothermic: products higher, arrow up labelled "energy absorbed".

BOND BREAKING AND MAKING

Breaking bonds requires energy (endothermic).
Making bonds releases energy (exothermic).
The overall energy change = energy needed to break bonds − energy released making bonds:
If more energy is released making bonds than used breaking them → exothermic (this is the case for most reactions).
If more is needed to break bonds → endothermic.
KEY POINT: bond breaking is ALWAYS endothermic; bond making is ALWAYS exothermic — the overall type depends on the balance.

MEASURING ENERGY CHANGES (calorimetry)

🔬 Practical — measuring the energy released by a fuel (e.g. burning food/alcohol): 1. Burn the fuel under a known volume/mass of water in a container. 2. Measure the temperature rise of the water. 3. Energy released (J) = mass of water (g) × 4.2 (J/g°C) × temperature rise (°C). 4. Divide by the mass of fuel burned to get energy per gram.
Limitations: heat lost to the surroundings → the measured value is lower than the true value; use a lid/insulation to improve accuracy.
🔬 Practical — temperature change in a reaction: mix reactants in an insulated cup (polystyrene) and record the temperature change; exothermic → rises; endothermic → falls.

✅ Quick check

Define exothermic and endothermic. (Exo = releases heat (temp rises); endo = absorbs heat (temp falls).)
Give two examples of each. (Exo: combustion, neutralisation. Endo: thermal decomposition, photosynthesis.)
Is bond breaking endothermic or exothermic? (Endothermic — requires energy.)
In an exothermic reaction, do products have more or less energy than reactants? (Less.)
How do you calculate the energy change from a calorimetry experiment? (E = mass of water × 4.2 × temperature rise.)

Ch 20 Rates of Reaction

Chapter 20: Rates of Reaction

The rate of a reaction is how fast reactants turn into products. Rates increase with higher temperature, higher concentration, higher pressure (gases), bigger surface area, and catalysts. All are explained by collision theory — more frequent and more energetic collisions.

MEASURING RATES

Rate = change in amount of reactant/product ÷ time. Ways to measure:
Mass loss (if a gas escapes — e.g. measuring CO₂ loss on a balance).
Volume of gas produced (gas syringe over time).
Turbidity/colour change (e.g. the disappearing cross with sodium thiosulfate + acid).
Change in pH/mass.
Rate is calculated from the gradient of a graph of quantity against time (steepest at the start, flattening as reactants run out).

THE FIVE FACTORS THAT SPEED UP REACTIONS

Temperature ↑ — particles have more kinetic energy → move faster → more collisions per second AND a greater proportion of collisions have enough energy (activation energy) to react.
Concentration ↑ (solutions) — more particles in the same volume → more frequent collisions.
Pressure ↑ (gases) — particles squeezed closer → more frequent collisions.
Surface area ↑ (solids crushed/powdered) — more particles exposed → more frequent collisions (powder reacts much faster than lumps).
Catalyst — provides an alternative reaction pathway with a lower activation energy → more collisions are successful (the catalyst is not used up).

COLLISION THEORY (the explanation you must give)

For a reaction to happen, particles must: 1. Collide with each other, AND 2. Collide with enough energy (at least the activation energy) and the right orientation.
Activation energy — the minimum energy particles need to react.
Any factor that increases the frequency of collisions or the proportion of successful collisions increases the rate.

CATALYSTS

Catalyst — a substance that speeds up a reaction without being used up (chemically unchanged at the end).
Catalysts are specific (each reaction has its own).
Enzymes are biological catalysts.
Examples: manganese dioxide speeds up the decomposition of hydrogen peroxide; iron in the Haber process; platinum/rhodium in catalytic converters; nickel in hydrogenation of oils.
Benefits: lower temperature/pressure needed → saves energy and money in industry.

PRACTICALS YOU SHOULD KNOW

🔬 Marble chips + acid: measure gas volume or mass loss; repeat with different surface areas/concentrations/temperatures.
🔬 Sodium thiosulfate + hydrochloric acid: time how long the "disappearing cross" takes; faster (higher concentration/temperature) = quicker.
🔬 Hydrogen peroxide + manganese dioxide (catalyst): measure oxygen volume over time; the catalyst makes bubbles appear faster.

✅ Quick check

Name the five factors that affect reaction rate. (Temperature, concentration, pressure, surface area, catalyst.)
State collision theory. (Particles must collide with at least the activation energy to react.)
Why does increasing temperature speed up a reaction? (Faster particles → more collisions + more successful (energetic) collisions.)
What is a catalyst? (Speeds up a reaction without being used up.)
Why does powdered solid react faster than lumps? (Bigger surface area → more frequent collisions.)

Ch 21 Reversible Reactions and Equilibria

Chapter 21: Reversible Reactions and Equilibria

Some reactions are reversible — they can go both ways. In a closed system they reach equilibrium, where the forward and backward reactions happen at the same rate. Le Chatelier's principle predicts how changing temperature, pressure or concentration shifts the equilibrium.

REVERSIBLE REACTIONS

Reversible reaction — a reaction that can go in both directions (shown with ⇌).
Example: hydrated copper sulfate ⇌ anhydrous copper sulfate + water (blue ⇌ white).
Example: ammonium chloride ⇌ ammonia + hydrogen chloride (heating drives it one way, cooling the other).
Reversible reactions are usually exothermic in one direction and endothermic in the other.

EQUILIBRIUM

Dynamic equilibrium — in a closed system (nothing escapes), the forward and backward reactions happen at the same rate, so the amounts of reactants and products stay constant.
KEY POINT: equilibrium is NOT "everything stopped" — both reactions continue; it is dynamic (the concentrations are constant).
Equilibrium is only reached in a closed system.

LE CHATELIER'S PRINCIPLE

If a system at equilibrium is disturbed, it shifts to counteract the change (to reduce the disturbance).
Used by industry to choose conditions for maximum yield.

1. CHANGING CONCENTRATION

Increasing the concentration of a reactant → equilibrium shifts to the right (uses it up) → more product.
Decreasing a product's concentration (removing it) → shifts right (replaces it).
Increasing a product → shifts left.

2. CHANGING PRESSURE (gases only)

Increasing pressure → equilibrium shifts toward the side with fewer gas molecules (reduces pressure).
Decreasing pressure → shifts toward the side with more gas molecules.
If both sides have the same number of gas molecules, pressure has no effect.

3. CHANGING TEMPERATURE

Increasing temperature → equilibrium shifts in the endothermic direction (absorbs the extra heat).
Decreasing temperature → shifts in the exothermic direction.
(Catalysts speed up both directions equally — they do NOT shift the equilibrium, just help it be reached faster.)

THE HABER PROCESS (making ammonia — the classic equilibrium question)

Reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) — exothermic forward reaction (4 gas molecules → 2).
Conditions chosen (compromise):
Temperature ~450 °C: lower temperature would give more ammonia (exothermic), but too slow — 450 °C is a compromise between yield and speed.
Pressure ~200 atmospheres (high): high pressure favours the side with fewer molecules (2 vs 4) → more ammonia. (Very high pressure is too expensive/dangerous.)
Iron catalyst: speeds up the reaction without affecting the yield.
Ammonia is continuously removed (shifts equilibrium right — see concentration rule).
📊 Exam tip: always explain each condition using Le Chatelier's principle AND the cost/speed compromise.

✅ Quick check

What is a reversible reaction? (A reaction that can go both ways — shown with ⇌.)
What is dynamic equilibrium? (Forward and backward reactions at the same rate in a closed system — concentrations constant.)
State Le Chatelier's principle. (A system at equilibrium shifts to counteract any change.)
In N₂ + 3H₂ ⇌ 2NH₃, what happens if pressure is increased? (Equilibrium shifts right (fewer gas molecules) → more ammonia.)
Why is the Haber process run at ~450 °C and 200 atm with an iron catalyst? (Compromise: enough speed/yield, high pressure favours product, catalyst speeds it up.)

UNIT 4 · Organic Chemistry

Ch 22 Introduction to Organic Chemistry

Chapter 22: Introduction to Organic Chemistry

Organic chemistry is the chemistry of carbon compounds. This chapter introduces the homologous series, the first series (alkanes), and the ideas of saturated/unsaturated, isomerism, and functional groups — the foundation for the chapters on crude oil and its products.

WHY CARBON?

Carbon forms four covalent bonds and can join to itself in chains and rings — this is why there are millions of organic compounds.

HOMOLOGOUS SERIES

Homologous series — a family of organic compounds with:
The same functional group (the reactive part).
The same general formula.
Similar chemical properties.
Each member differs from the next by CH₂.
The series you must know: alkanes, alkenes, alcohols, carboxylic acids, esters (and crude oil products).

THE ALKANES (the first series)

Alkanes — saturated hydrocarbons (only single C–C bonds); general formula CₙH₂ₙ₊₂.
Hydrocarbon — a compound containing ONLY hydrogen and carbon.
First four alkanes (memorise): | Name | Formula | Structure | |---|---|---| | Methane | CH₄ | | | Ethane | C₂H₆ | CH₃CH₃ | | Propane | C₃H₈ | CH₃CH₂CH₃ | | Butane | C₄H₁₀ | CH₃CH₂CH₂CH₃ |
Saturated — contains only single bonds (no more hydrogen can be added).
Alkanes are quite unreactive (burn in oxygen; react with halogens in UV light).

ISOMERS

Isomers — compounds with the same molecular formula but different structural formulae (atoms connected in a different order).
Example: butane C₄H₁₀ has two isomers — straight-chain butane and branched methylpropane (different boiling points).

FUNCTIONAL GROUPS

Functional group — the atom/group in a molecule that gives it its characteristic reactions:
Alkanes: none (single bonds only).
Alkenes: C=C double bond.
Alcohols: –OH.
Carboxylic acids: –COOH.
Esters: –COO–.
KEY POINT: the functional group determines the family's chemistry (and naming: -ane, -ene, -ol, -oic acid, -oate).

HOW TO DRAW AND NAME ORGANIC COMPOUNDS

Types of formula:
Molecular formula — the actual numbers of atoms (C₂H₆).
Structural formula — shows how atoms are joined, written in a line (CH₃CH₃ for ethane; CH₃CH₂CH₂CH₃ for butane).
Displayed (graphic) formula — shows every bond between every atom (drawn out).
Naming (the code): the prefix tells the chain length (meth- = 1, eth- = 2, prop- = 3, but- = 4); the suffix tells the family (-ane, -ene, -ol, -oic acid, -oate).
Branched chains: number the longest chain so the branches get the lowest numbers; name the branch with a prefix (methyl = 1C branch, ethyl = 2C branch) and its position. E.g. 2-methylpropane (butane's isomer), 2-methylbutane.
💡 Exam tip: when drawing structural formulae, check every carbon makes four bonds.

✅ Quick check

What is a hydrocarbon? (A compound of only hydrogen and carbon.)
What is a homologous series? (A family with the same functional group, same general formula, similar properties, differing by CH₂.)
Give the general formula of alkanes. (CₙH₂ₙ₊₂.)
Name the first four alkanes. (Methane, ethane, propane, butane.)
What are isomers? (Same molecular formula, different structural formula.)

Ch 23 Crude Oil

Chapter 23: Crude Oil

Crude oil is a fossil fuel — a mixture of many hydrocarbons (mainly alkanes). It is separated by fractional distillation into fractions with different boiling points, each with its uses. This chapter covers the process, the fractions, and cracking (making smaller, more useful molecules).

WHAT IS CRUDE OIL?

Crude oil — a dark, thick liquid formed from the remains of tiny marine organisms over millions of years (a non-renewable fossil fuel).
It is a mixture of hydrocarbons (mainly alkanes of different chain lengths).
Because it is a mixture, it has no single boiling point — the components can be separated by fractional distillation.

FRACTIONAL DISTILLATION

The crude oil is vaporised and fed into a fractionating column (hot at the bottom, cool at the top).
How it works:
Short-chain molecules have low boiling points → stay gaseous → rise to the top of the column.
Long-chain molecules have high boiling points → condense early → collect near the bottom.
Each fraction collects at the level where its boiling point matches the temperature.
KEY POINT: fraction = a group of hydrocarbons with similar boiling points.

THE FRACTIONS (bottom → top of the column)

Fraction	Chain length	Boiling point	Uses
Refinery gases (e.g. butane)	1–4	Very low	Bottled gas (camping gas), heating
Petrol (gasoline)	~5–9	Low	Fuel for cars
Naphtha	~6–10	Low	Making chemicals/plastics
Kerosene (paraffin)	~10–16	Medium	Jet fuel
Diesel (gas oil)	~14–20	Medium-high	Diesel engines, trains
Fuel oil	~20+	High	Ships, power stations, heating
Bitumen	Very long	Very high	Surfacing roads, roofing
- Trends (memorise): the LONGER the chain → higher boiling point, thicker (more viscous), less volatile, darker, harder to ignite.

CRACKING

Why crack? Demand is high for short-chain molecules (petrol) but crude oil contains too many long chains — cracking converts long chains into smaller, more useful ones.
Cracking — breaking down long-chain hydrocarbons into smaller, more useful molecules (a shorter alkane + an alkene, which is useful for making plastics).
Conditions: high temperature (~400–700 °C) + a catalyst (catalytic cracking), or steam (steam cracking).
Example: C₁₀H₂₂ → C₈H₁₈ + C₂H₄ (decane → octane + ethene)
Products: a smaller alkane (fuel) and an alkene (e.g. ethene — used to make ethanol and polymers like poly(ethene)).

COMBUSTION OF HYDROCARBONS (fuel use)

Complete combustion (plenty of oxygen): hydrocarbon + O₂ → CO₂ + H₂O (clean flame, releases lots of energy).
Incomplete combustion (not enough oxygen): produces CO (toxic) or carbon (soot) + water.
Burning fossil fuels adds CO₂ (global warming) and other pollutants (see Ch 13).

✅ Quick check

What is crude oil? (A mixture of hydrocarbons, mainly alkanes — a non-renewable fossil fuel.)
How is crude oil separated? (Fractional distillation — by boiling point in a fractionating column.)
Name the fractions from top to bottom. (Refinery gases, petrol, naphtha, kerosene, diesel, fuel oil, bitumen.)
Why is cracking needed? (Demand for short chains (petrol) exceeds supply — crack long chains into short alkane + alkene.)
Write an example cracking equation. (C₁₀H₂₂ → C₈H₁₈ + C₂H₄.)

Ch 24 Alkanes

Chapter 24: Alkanes

Alkanes are saturated hydrocarbons (only single C–C bonds). They are found in crude oil and natural gas and are used mainly as fuels. This chapter covers their properties, combustion, and their reaction with halogens (substitution).

THE ALKANES — RECAP

General formula: CₙH₂ₙ₊₂.
Saturated — every carbon forms four single bonds (no double bonds) → they are quite unreactive.
Members: methane (CH₄), ethane (C₂H₆), propane (C₃H₈), butane (C₄H₁₀)...
Physical properties trend (longer chain): higher boiling point, less volatile, more viscous, less flammable.

COMBUSTION (burning — the most important reaction)

Alkanes are excellent fuels — they burn in oxygen releasing energy.
Complete combustion (enough oxygen): alkane + oxygen → carbon dioxide + water CH₄ + 2O₂ → CO₂ + 2H₂O Blue, clean flame; releases lots of energy.
Incomplete combustion (limited oxygen):
Produces carbon monoxide (CO) — toxic (binds to haemoglobin), and/or soot (carbon).
Example: 2CH₄ + 3O₂ → 2CO + 4H₂O.
Dangerous in poorly ventilated heaters/gas appliances (CO poisoning — silent killer; need CO alarms).
💡 Exam tip: "explain why a yellow flame and black soot appear" → incomplete combustion, limited oxygen.

REACTION WITH HALOGENS (SUBSTITUTION)

Alkanes react with chlorine/bromine in the presence of UV light — a substitution reaction (a hydrogen atom is replaced by a halogen atom). CH₄ + Cl₂ →(UV light)→ CH₃Cl + HCl
The UV light provides the energy to break the bonds (homolytic fission forms radicals).
KEY POINT: this is how halogenated organic compounds (e.g. CFCs, solvents) are made.

USES OF ALKANES

Fuels: methane (natural gas — heating/cooking), propane/butane (bottled gas), petrol/diesel (from crude oil fractions).
Feedstock for making other chemicals (cracking → alkenes; chlorination → solvents).

✅ Quick check

What does "saturated" mean? (Only single bonds between carbons — no double bonds.)
Write the complete combustion of methane. (CH₄ + 2O₂ → CO₂ + 2H₂O.)
What is produced by incomplete combustion? (Carbon monoxide and/or carbon (soot).)
Why is carbon monoxide dangerous? (It is toxic — binds to haemoglobin, reducing oxygen transport.)
What conditions are needed for alkanes to react with chlorine? (UV light — substitution reaction.)

Ch 25 Alkenes

Chapter 25: Alkenes

Alkenes are unsaturated hydrocarbons containing a C=C double bond — their functional group. They are much more reactive than alkanes and are the starting materials for making polymers and many other chemicals.

THE ALKENES

General formula: CₙH₂ₙ (one less hydrogen than alkanes — because of the double bond).
Unsaturated — contain a C=C double bond (they can react by opening the double bond and adding more atoms).
First members: ethene (C₂H₄) and propene (C₃H₆).
Testing for unsaturation (KEY practical):
Add bromine water (orange/brown) and shake:
- Alkene (unsaturated): the bromine water is decolourised (turns colourless) — the bromine adds across the double bond (addition reaction).
- Alkane (saturated): stays orange/brown.
💡 Exam tip: decolourising bromine water is THE test for a double bond.

WHY ALKENES ARE MORE REACTIVE

The C=C double bond is an area of high electron density — it attracts other reactants (electrophiles) and can open up so atoms add to each carbon. Alkanes (single bonds) have no such reactive site.

ADDITION REACTIONS OF ALKENES

Alkenes undergo addition reactions (the double bond opens, atoms add — no atoms are lost). 1. With hydrogen (hydrogenation): alkene + H₂ → alkane (needs a nickel catalyst and heat). Ethene + H₂ → ethane: C₂H₄ + H₂ → C₂H₆. (Used to harden oils into margarine.) 2. With halogens (e.g. bromine): ethene + bromine → dibromoethane: C₂H₄ + Br₂ → C₂H₄Br₂ (this is the bromine water test reaction). 3. With hydrogen halides: ethene + HCl → chloroethane: C₂H₄ + HCl → C₂H₅Cl. 4. With steam (hydration): ethene + steam → ethanol: C₂H₄ + H₂O → C₂H₅OH (needs a catalyst — phosphoric acid, high temperature and pressure). This is an industrial way to make ethanol.

POLYMERISATION (the big one)

Polymerisation — many small molecules (monomers) join together to form a very long molecule (polymer).
Alkenes polymerise by addition polymerisation (the double bond opens and the monomers link in a chain).
Example — making poly(ethene): n C₂H₄ → poly(ethene) [–CH₂–CH₂–]ₙ (high pressure + catalyst).
KEY POINT: the polymer has the same atoms as the monomers (nothing is lost) — addition polymerisation.
Poly(propene), poly(chloroethene) (PVC), poly(tetrafluoroethene) (PTFE) are other examples.
Uses: plastic bags and bottles (poly(ethene)), pipes and window frames (PVC), non-stick coatings (PTFE).
💡 Exam tip: be able to draw the monomer from the polymer repeat unit and vice versa.

✅ Quick check

What is the functional group of alkenes? (The C=C double bond.)
What is the test for an alkene? (Decolourises bromine water — addition reaction.)
Write the hydrogenation of ethene. (C₂H₄ + H₂ → C₂H₆, nickel catalyst.)
What is addition polymerisation? (Monomers with double bonds join into a long chain — nothing is lost.)
Write the equation for making ethanol from ethene. (C₂H₄ + H₂O → C₂H₅OH, with a catalyst.)

Ch 26 Alcohols

Chapter 26: Alcohols

Alcohols have the functional group –OH and are made from alkenes (hydration) or by fermentation. They are used as fuels, solvents and in drinks. This chapter covers ethanol's structure, properties, reactions (combustion, oxidation) and the two ways of making it.

THE ALCOHOLS

Functional group: –OH (hydroxyl group).
General formula: CₙH₂ₙ₊₁OH (or CₙH₂ₙ₊₂O).
Ethanol (C₂H₅OH) is the most important — found in alcoholic drinks.
Names end in -ol: methanol (CH₃OH), ethanol (C₂H₅OH), propanol (C₃H₇OH).

PROPERTIES OF ETHANOL

Colourless liquid at room temperature.
Miscible with water (mixes fully — the –OH group forms hydrogen bonds with water).
Volatile (low boiling point — evaporates easily).
Flammable — burns well.

REACTIONS OF ETHANOL

Combustion: ethanol burns in oxygen → carbon dioxide + water (releases energy). C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O Used as a biofuel (ethanol from fermentation mixed with petrol — "gasohol").
Oxidation (to ethanoic acid): ethanol is oxidised by acidified potassium dichromate (K₂Cr₂O₇) — the orange dichromate turns green, and ethanol becomes ethanoic acid (vinegar's acid). - This is the basis of the breathalyser test.
Dehydration: removing water (with concentrated sulfuric acid catalyst) → ethene: C₂H₅OH → C₂H₄ + H₂O.

MAKING ETHANOL — TWO METHODS (compare them!)

Method 1: HYDRATION OF ETHENE (industrial)

Ethene (from cracking) + steam, with a phosphoric acid catalyst, high temperature and pressure. C₂H₄ + H₂O → C₂H₅OH
Advantages: continuous process, pure ethanol, no waste.
Disadvantages: ethene comes from non-renewable crude oil.

Method 2: FERMENTATION (biological)

Sugar (glucose) + yeast, warm (~30–35 °C), no oxygen: C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂
Advantages: uses renewable plant material (sugar cane, maize); cheap.
Disadvantages: slow, batch process; produces a dilute solution (~15% max — yeast dies in higher alcohol) so it must be distilled to concentrate; produces CO₂.

✅ Quick check

What is the functional group of alcohols? (–OH.)
Write the combustion of ethanol. (C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O.)
What happens when ethanol is oxidised? (It becomes ethanoic acid; orange dichromate turns green.)
Write the fermentation equation. (C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂, with yeast.)
Give one advantage and one disadvantage of each method of making ethanol. (Hydration: pure/continuous but non-renewable oil. Fermentation: renewable but slow, dilute, needs distillation.)

Ch 27 Carboxylic Acids

Chapter 27: Carboxylic Acids

Carboxylic acids contain the functional group –COOH. They are weak acids found in vinegar (ethanoic acid) and many foods. This chapter covers their structure, properties, reactions, and how they're made (oxidation of alcohols).

THE CARBOXYLIC ACIDS

Functional group: –COOH (carboxyl group).
Names end in -oic acid: methanoic acid (HCOOH), ethanoic acid (CH₃COOH) — the acid in vinegar.
General formula: CₙH₂ₙ₊₁COOH.

PROPERTIES

Weak acids — partially ionise in water (they turn litmus red, have a pH around 3–5).
React like other acids:
With metals → salt + hydrogen.
With bases/alkalis → salt + water.
With carbonates → salt + water + carbon dioxide (fizzes — limewater test).
The salts are called ethanoates, methanoates, etc. (e.g. sodium ethanoate).
KEY POINT: carboxylic acids react more slowly/weakly than strong mineral acids (HCl, H₂SO₄) because they only partially ionise.

MAKING ETHANOIC ACID

Oxidation of ethanol (see Ch 26): ethanol + oxidising agent (acidified potassium dichromate) → ethanoic acid. This is how vinegar is made — bacteria oxidise ethanol in wine/cider to ethanoic acid.

REACTION WITH ALCOHOLS — MAKING ESTERS (link to Ch 28)

Carboxylic acid + alcohol ⇌ ester + water (an esterification reaction, needs a strong acid catalyst — concentrated sulfuric acid): ethanoic acid + ethanol ⇌ ethyl ethanoate + water CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O
Esters have fruity smells and are used as flavourings and perfumes (see Ch 28).

USES OF CARBOXYLIC ACIDS

Ethanoic acid: vinegar (food preservative and flavouring), making esters, descaling.
Methanoic acid: in some insect stings (formic acid), leather tanning.

✅ Quick check

What is the functional group of carboxylic acids? (–COOH.)
Give the formula of ethanoic acid. (CH₃COOH.)
Why are carboxylic acids described as weak acids? (They only partially ionise in water.)
What happens when a carboxylic acid reacts with a carbonate? (Salt + water + carbon dioxide — fizzes.)
How is ethanoic acid made? (By oxidising ethanol — e.g. with acidified potassium dichromate, or bacteria in vinegar-making.)

Ch 28 Esters

Chapter 28: Esters

Esters are formed from a carboxylic acid + an alcohol (esterification). They have sweet, fruity smells and are used in flavourings, perfumes and as solvents. This chapter covers their formation, properties and uses.

WHAT ARE ESTERS?

Ester — a compound formed by the reaction of a carboxylic acid with an alcohol.
Functional group: –COO– (the ester link).
Names end in -oate: the alcohol part first, then the acid part — e.g. ethyl ethanoate (from ethanol + ethanoic acid).

FORMATION — ESTERIFICATION

carboxylic acid + alcohol ⇌ ester + water

Example: CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O ethanoic acid + ethanol ⇌ ethyl ethanoate + water
Conditions: warm with a strong acid catalyst (concentrated sulfuric acid).
The reaction is reversible (⇌).
Other examples: methyl ethanoate (methanol + ethanoic acid), ethyl propanoate (ethanol + propanoic acid), butyl ethanoate (butanol + ethanoic acid — smells of bananas).

PROPERTIES OF ESTERS

Sweet, fruity smells (e.g. ethyl ethanoate smells like pear drops/glue).
Volatile liquids (evaporate easily — that's why you can smell them).
Insoluble in water (but soluble in organic solvents).
Lower boiling points than the parent acid/alcohol.

USES OF ESTERS

Flavourings — synthetic fruit flavours in sweets, drinks and ice cream.
Perfumes and cosmetics — their smells evaporate and linger.
Solvents — e.g. in nail polish remover, glues, paints (ethyl ethanoate is a common solvent).
Plasticisers and making polyesters (see Ch 29).

💡 Exam tip

To name an ester: alkyl group from the alcohol (ethyl, methyl...) + -oate from the acid (ethanoate, propanoate...).
To draw it: the alcohol part (–O–) links the two halves.

MAKING AND TESTING ESTERS (practical)

🔬 Practical — making ethyl ethanoate: warm a mixture of ethanoic acid, ethanol and a few drops of concentrated sulfuric acid (catalyst) in a water bath; pour into water — the ester separates as an oily layer with a sweet, fruity smell (smell by wafting, never sniff directly).
Other esters: methyl ethanoate (smells of pear drops), butyl ethanoate (bananas), propyl ethanoate (pears). The different smells come from the different alcohol parts.

✅ Quick check

What is an ester? (A compound formed from a carboxylic acid and an alcohol.)
Write the equation for making ethyl ethanoate. (CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O.)
What conditions are needed? (Warm with concentrated sulfuric acid catalyst.)
Give two uses of esters. (Flavourings, perfumes, solvents.)
Why do esters smell strong? (They are volatile — evaporate easily.)

Ch 29 Synthetic Polymers

Chapter 29: Synthetic Polymers

Polymers are very long molecules made of repeating units (monomers). This chapter covers addition polymers (from alkenes) and condensation polymers (polyesters, polyamides), the properties of plastics, and the environmental issues of plastics and recycling.

ADDITION POLYMERS

Made by addition polymerisation: alkene monomers open their double bonds and join in a chain — no other product is formed (nothing lost).
Example — poly(ethene): n C₂H₄ → [–CH₂–CH₂–]ₙ (ethene → poly(ethene)).
The polymer's name comes from the monomer: ethene → poly(ethene); propene → poly(propene); chloroethene → poly(chloroethene) (PVC).
Conditions: high pressure + catalyst (for poly(ethene)).
KEY SKILL: from a monomer, draw the repeat unit (the part in brackets); from a repeat unit, identify the monomer (re-add the double bond).

CONDENSATION POLYMERS (Chemistry only)

Made by condensation polymerisation: monomers join with the loss of a small molecule (usually water) at each link.
The monomers have two functional groups each.
Polyesters: made from a dicarboxylic acid + a diol — joined by ester links (–COO–); water is released.
Polyamides (e.g. nylon, Kevlar): made from a dicarboxylic acid + a diamine — joined by amide links (–CONH–); water released.
KEY DIFFERENCE: addition polymers lose nothing; condensation polymers lose water (or another small molecule) at each join.

PROPERTIES OF PLASTICS

Thermoplastics — soften and melt when heated (can be reshaped and recycled). The polymer chains are held by weak forces — e.g. poly(ethene), nylon, PVC.
Thermosetting plastics (thermosets) — do NOT melt when heated (they set permanently); strong cross-links between chains — e.g. Bakelite, melamine, epoxy resins. Used for electrical fittings, kettles, pans.
💡 Exam tip: know the difference — thermoplastics can be remoulded; thermosets cannot (cross-linked).

PROBLEMS WITH PLASTICS

Made from non-renewable crude oil.
Non-biodegradable — most do not rot; they accumulate in landfill and the oceans (plastic pollution harms wildlife).
Burning them can release toxic fumes (e.g. HCl from PVC).
Recycling plastics saves oil, reduces landfill and pollution — but plastics must be sorted (different types melt at different temperatures) and some are hard to recycle (thermosets).

SOLUTIONS

Recycle (sort and melt down thermoplastics).
Biodegradable polymers (from plant material like starch/corn — break down naturally).
Reuse and reduce (less packaging).
Burn for energy recovery (waste-to-energy, carefully filtered).

✅ Quick check

What is an addition polymer? (Alkene monomers join by opening the double bond — nothing else is formed.)
What is a condensation polymer? (Monomers join with the loss of a small molecule — usually water — e.g. polyesters, polyamides.)
Distinguish thermoplastics and thermosets. (Thermoplastics melt and can be reshaped; thermosets are cross-linked and cannot.)
Give two environmental problems of plastics. (Non-biodegradable — landfill/ocean pollution; made from non-renewable oil; toxic fumes when burned.)
Name two ways to reduce plastic waste. (Recycling, biodegradable polymers, reduce/reuse.)

📖 Glossary of Key Terms

Glossary of Key Terms

Extracted from the back of the textbook (OCR — some entries may need light editing).

328 GLOSSARY GLOSSARY absorb take in (a gas or a liquid) anode the positive electrode in electrolysis, which (Ht) in solution (Arrhenius theory) or as a proton donor acid a substance that acts as a source of hydrogen ions attracts negative anions (Brønsted-Lowry theory) pattern of the others anomalous result a result that does not fit in with the acid rain rain which has a pH of less than about 5.6 react with sulfur dioxide to produce sulfuric acid, or with It is caused when water and oxygen in the atmosphere aqueous something dissolved in water array an ordered arrangement of things various oxides of nitrogen, NO,, to give nitric acid activation energy the minimum amount of energy been burned ash the soft powder that remains after something has reaction required for a collision to be successful, i.e. to result in a recognised as that element atom the smallest piece of an element that can still be adds to another without taking anything away, to form addition a chemical reaction in which one molecule atomic number the number of protons in an atom with halogens, the halogen atoms add onto the alkene a single product. For example, when alkenes react 12C (6.02 x 1023 mol-1) Avogadro constant the number of 12C atoms in 12g of molecule addition polymerisation polymerisation of monomers in front of formulae so that the same number of atoms of balancing the equation a process of putting coefficients containing a carbon-carbon double bond. A large number anything else being formed of monomer molecules add onto each other without each type is on both side of an equation iron with paint, oil, grease or plastic, so that oxygen/water barrier method a method of rust prevention by coating alcohols all contain an -OH functional group attached to a a homologous series of compounds which cannot reach the iron/steel hydrocarbon chain with the hydrogen ions in them. They are usually metal base a substance that neutralises acids by combining alkali metals group 1 elements including lithium, Note, hydrogen is not an alkali metal sodium, potassium, rubidium, caesium and francium oxides, hydroxides or ammonia. A soluble base is called hydroxide ions (OH) in solution (Arrhenius theory) or as a an alkali, and it is a substance that acts as a source of which all the carbons are joined to each other with single alkanes a homologous series of similar hydrocarbons in proton acceptor (Brønsted-Lowry theory) covalent bonds. These are saturated compounds with the with together batch a group of things that are produced or are dealt general formula C,H2n+2 alkenes a homologous series of hydrocarbons which as sugar cane or corn biofuel a fuel that is made from biological sources, such contain a carbon-carbon double bond. These are unsaturated compounds with the general formula C,H2n allotropes different forms of the same element, for biopolyester a polyester which is biodegradable allotropes of carbon example diamond, graphite and Ceo fullerene are three whiter, or to sterilise something bleach a chemical used to make something paler or carbon. For example, brass is an alloy of copper and zinc, alloy a mixture of a metal with, usually, other metals or boiling the change of state from a liquid to a gas. It occurs at the boiling point and steel is an alloy of iron and carbon bond energy the amount of energy required to break amphoteric substances that can react with both acids amount of energy released when 1 mole of covalent 1 mole of covalent bonds in gaseous molecules, or the and bases to form salts anaerobic in the absence of air bonds are formed in gaseous molecules anhydrous without water an object branch something that grows out from the main part of anion a negative ion, formed by atoms gaining electrons bubbling producing bubbles

GLOSSARY 329 is passed through it. Or, the wider part of some plastic bulb a glass object that produces light when electricity a small molecule such as water or hydrogen chloride is monomers in which each time two monomers combine, condensation polymerisation polymerisation of filled pipettes, which is squeezed to allow the pipette to be removed a chemical reaction calorimetry measuring the heat given out or taken in by two molecules combine to form a larger molecule with condensation reaction a chemical reaction in which which all contain a -COOH functional group attached to a carboxylic acids a homologous series of compounds elimination of a small molecule such as water hydrocarbon chain corrosive a substance which can damage living tissue catalyst substance that speeds up a chemical reaction covalent bonding strong electrostatic force of attraction energy. Catalysts are not used up and remain chemically by providing an alternative pathway of lower activation the shared pair of electrons between the nuclei of the atoms making up the bond and unchanged at the end of the reaction converted to alkenes and shorter-chain alkanes. It cracking a process in which long-chain alkanes are catalytic converter a device used in cars to convert oxides of nitrogen and carbon monoxide into harmless temperature of 600-700°C is carried out using silica or alumina as catalyst at a palladium and rhodium as catalysts nitrogen gas and carbon dioxide. It uses platinum, when their soft tissue was gradually changed by high crude oil formed from the remains of living organisms cathode the negative electrode in electrolysis, which attracts positive cations mixture of hydrocarbons temperatures and pressures into a thick, black oil. It is a cation a positive ion, formed by atoms losing electrons solid) is obtained from its solvent crystallisation a process in which a solute (soluble reactions chemical means methods which involve chemical damp slightly wet chromatography showing the separation of different chromatogram the adsorbent paper from paper a compound is broken down into its elements or simpler decomposition reaction a chemical reaction in which coloured substances compounds clump a small group or cluster dehydration removal of water balanced chemical equation coefficient number written in front of formulae in a attached to particular atoms or pairs of atoms, but are delocalised electrons electrons that are no longer coil continuous series of circular rings into which free to move throughout the whole structure something such as wire or rope has been wound or twisted structural changes, usually caused by changes in denatured the loss of function of enzymes due to collision theory states that for a reaction to occur, the collision needs to have sufficient energy and the correct reactant particles must collide with each other and the temperature or pH orientation experiment dependent variable what you measure in an reacts with oxygen (burns) to form products and heat combustion a chemical reaction in which a substance solid deposition the change of state directly from a gas to a competition reaction displacement reaction between a more reactive metal and the oxide of a less reactive metal diatomic a molecule that contains two atoms complete combustion occurs when a hydrocarbon area of high concentration to an area of low diffusion the random movement of particles from an burns in sufficient oxygen (burns) and forms carbon dioxide and water as products concentration elements chemically combine The elements cannot be compound a substance that forms when two or more react to form a larger molecule dimerise forming a dimer - two identical molecules separated by physical means dip put something into a liquid and lift out again discharged when an ion loses its charge by losing or It occurs at the condensation point condensation the change of state from a gas to a liquid. gaining electrons

330 GLOSSARY more reactive element replaces a less reactive one in its displacement reaction a chemical reaction in which a from the surroundings endothermic reactions in which heat energy is taken in compound a molecule as individual lines. Each line represents a pair displayed formula a