If you go on to the full A Level, these three papers make up 50% of it. Papers 4 and 5, in the A Level tab, make up the other 50%.

Topics

How the marks work

Chemistry mark schemes list the marking points separately. Each point (labelled M1, M2 and so on) is worth one mark. The same words come up again and again in them:

Command words

The Data Booklet

In every theory paper you get a Data Booklet with a Periodic Table and tables of data. You don't need to memorise these values:

The booklet also lists ionisation energies, bond energies, standard electrode potentials, atomic and ionic radii, and the infrared and NMR absorption tables. Values in different books differ slightly, so in the exam always use the values printed in the booklet or in the question. The questions on this site give any data you need.

Almost everything in chemistry comes down to electrons: where they are, how strongly the nucleus holds them, and how easily they are lost or shared. This topic builds that picture, from the particles in the atom to orbitals and ionisation energies.

Inside the atom

An atom is mostly empty space. At its centre is a very small, dense nucleus that contains the protons and neutrons. The nucleus is about 10⁻¹⁵ m across, while the whole atom is about 10⁻¹⁰ m across, roughly 100 000 times larger. The electrons are found in shells (energy levels) in the space around the nucleus.

Distribution of mass and charge: nearly all the mass of an atom is in the nucleus, which is positively charged. The negative charge is spread through the much larger volume around it. An atom is neutral because it has equal numbers of protons and electrons.

Beams of particles in an electric field

If beams of protons, neutrons and electrons travel at the same speed between two charged plates:

• Protons are deflected towards the negative plate.
• Electrons are deflected towards the positive plate, and through a much larger angle because their mass is about 1/1836 of a proton's, but the size of their charge is the same.
• Neutrons are not deflected because they have no charge.

Counting protons, neutrons and electrons

Proton number (atomic number), Z

The number of protons in the nucleus. It identifies the element.

Nucleon number (mass number), A

The total number of protons and neutrons in the nucleus.

A nuclide is written 3717Cl, with the nucleon number at the top and the proton number at the bottom. So:

• number of protons = Z
• number of neutrons = A − Z
• number of electrons = Z − (charge on the ion). A 2+ ion has lost 2 electrons; a 2− ion has gained 2.

Atomic radius and ionic radius

The atomic radius is usually measured as half the distance between the nuclei of two bonded atoms of the same element. Two trends matter:

Ions compared with atoms

• Positive ions are smaller than their atoms. A metal atom such as Na loses its whole outer shell, and the remaining electrons are pulled in more strongly because there are now more protons than electrons.
• Negative ions are larger than their atoms. The extra electron(s) increase repulsion between electrons, while the nuclear charge stays the same, so the electron cloud spreads out.

An isoelectronic series (ions with the same number of electrons) shows the effect of nuclear charge clearly. All of these ions have the configuration 1s² 2s² 2p⁶:

Ten electrons are held by more and more protons, so the radius falls steadily.

Across Period 3

The cations Na⁺ to Si⁴⁺ all have 2 shells of electrons and rising nuclear charge, so they get smaller. The anions P³⁻ to Cl⁻ have 3 full shells, which is why there is a jump in size between Si⁴⁺ and P³⁻. They then get smaller too, as nuclear charge rises.

Isotopes

Isotopes

Atoms of the same element with the same number of protons but different numbers of neutrons.

For example, chlorine has two isotopes, 3517Cl (18 neutrons) and 3717Cl (20 neutrons). Carbon has 126C, 136C and the radioactive 146C.

Shells, sub-shells and orbitals

Shell

A main energy level, numbered by the principal quantum number, n = 1, 2, 3, … The higher n is, the further the electrons are from the nucleus and the higher their energy.

Sub-shell

A group of orbitals of the same type within a shell: s, p, d or f.

Orbital

A region of space around the nucleus that can hold up to two electrons, with opposite spins.

Ground state

The lowest-energy arrangement of the electrons in an atom.

The order of filling

Three rules decide where the electrons go in the ground state:

1. Lowest energy first: electrons fill the sub-shell of lowest energy first, so 4s fills before 3d.
2. Two per orbital: an orbital holds at most two electrons, and they must have opposite spins (shown ↑↓).
3. Single first (Hund's rule): in a sub-shell with several orbitals of equal energy, electrons go into separate orbitals with parallel spins before they pair up. This is because two electrons in the same orbital repel each other more than two electrons in different orbitals.

Writing electronic configurations

The full configuration lists every sub-shell with the number of electrons as a superscript. The shorthand form starts with the previous noble gas in square brackets.

Cambridge writes 3d before 4s, grouping each shell together, even though 4s fills first. Either order is accepted.

Configurations of ions

• Negative ions: add electrons in the normal filling order. O²⁻ is 1s² 2s² 2p⁶.
• Positive ions of the d-block: remove the 4s electrons first, then 3d. Once the 3d orbitals contain electrons, 4s is the higher-energy, outer sub-shell.

Electrons in boxes

Each box is an orbital and each arrow is an electron. Arrows in the same box point in opposite directions. Use the explorer below to see the boxes for any element up to krypton, and for its ions.

Shapes of s and p orbitals

Orbitals in higher shells are the same shape but larger: a 2s orbital is bigger than a 1s orbital.

Free radicals

A free radical is a species with one or more unpaired electrons. It is shown with a dot, for example the chlorine atom Cl• and the methyl radical CH₃•. Nitrogen monoxide, NO, and nitrogen dioxide, NO₂, are also free radicals because each molecule has an odd number of electrons. Radicals are very reactive, as you will see in Topic 14.

Ionisation energy

First ionisation energy

The energy needed to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1+ ions.

Each successive ionisation removes one more electron from the ion formed in the step before. All ionisation energies are positive (endothermic), because energy is needed to pull an electron away from the attraction of the nucleus. Always include the (g) state symbols.

What decides the size of an ionisation energy?

Ionisation energy measures the attraction between the nucleus and the electron being removed. Four factors control it:

1. Nuclear charge: more protons attract the electron more strongly.
2. Distance from the nucleus (atomic or ionic radius): an electron further out is attracted less.
3. Shielding: inner shells and sub-shells repel the outer electron and reduce the attraction it feels.
4. Spin-pair repulsion: an electron sharing an orbital with another is repelled by it, so it is slightly easier to remove.

Trends in first ionisation energy

There are two small dips in each period, which give evidence for sub-shells:

The big drop from a noble gas to the next Group 1 metal (Ne → Na) happens because the outer electron of Na is in a new shell, further away and shielded by two full inner shells.

Successive ionisation energies

For any element, each successive ionisation energy is larger than the one before. After each electron is removed, the remaining electrons are attracted by the same number of protons, there is less repulsion between them, and the ion is smaller.

A big jump appears whenever the next electron has to come from a shell closer to the nucleus. Counting the electrons removed before the first big jump gives the number of outer-shell electrons, and so the group.

Questions

Questions 1 to 5 are multiple choice, like Paper 1. Questions 6 to 8 are structured, like Paper 2.

Stoichiometry is the arithmetic of chemistry: using formulas and balanced equations to work out how much of each substance reacts and forms. Almost every Paper 2 question and every titration in Paper 3 uses it, so practise until it is automatic.

Relative masses

Atoms are far too light to weigh in grams, so their masses are compared with a standard, the carbon-12 atom.

Unified atomic mass unit, u

One twelfth of the mass of one atom of carbon-12. (1 u = 1.66 × 10⁻²⁷ kg.)

Relative isotopic mass

The mass of one atom of an isotope compared with one twelfth of the mass of one atom of carbon-12.

Relative atomic mass, A_r

The weighted average mass of the atoms of an element (taking account of the abundance of its isotopes), compared with one twelfth of the mass of one atom of carbon-12.

Relative molecular mass, M_r

The weighted average mass of a molecule, compared with one twelfth of the mass of one atom of carbon-12. It is the sum of the A_r values of the atoms in the molecule.

Relative formula mass

The same idea for ionic compounds, which do not contain molecules, for example NaCl. It is often still written M_r.

Relative masses have no units, because they are ratios.

The abundances often come from a mass spectrum, where the height of each peak is proportional to the abundance of that isotope. See Topic 22.

The mole and the Avogadro constant

Mole

The amount of substance that contains 6.02 × 10²³ particles (atoms, molecules, ions or electrons). This number is the Avogadro constant, L.

Molar mass, M

The mass of one mole of a substance, in g mol⁻¹. It has the same numerical value as A_r or M_r.

Formulas

Ions you must know

The charge on a simple ion can be predicted from the group: Group 1 forms 1+, Group 2 forms 2+, Group 13 forms 3+, Group 15 forms 3−, Group 16 forms 2− and Group 17 forms 1−. Learn these compound ions and two common d-block ions:

For elements with more than one possible charge, a Roman numeral gives the oxidation number: iron(II) is Fe²⁺, iron(III) is Fe³⁺, copper(II) is Cu²⁺, and manganese(IV) oxide is MnO₂.

Writing formulas of ionic compounds

The total positive charge must equal the total negative charge. Use brackets when there is more than one compound ion.

Equations

A balanced equation has the same number of each type of atom, and the same total charge, on both sides. State symbols show the physical state: (s) solid, (l) liquid, (g) gas, (aq) dissolved in water.

An ionic equation shows only the species that change. Spectator ions, which are present but unchanged, are left out.

Empirical and molecular formulas

Empirical formula

The simplest whole-number ratio of the atoms of each element in a compound.

Molecular formula

The actual number of atoms of each element in one molecule of the compound.

For example, ethane C₂H₆ has the empirical formula CH₃. Glucose C₆H₁₂O₆ has the empirical formula CH₂O.

Hydrated salts

Hydrated

A crystalline compound that contains water molecules in its lattice, such as CuSO₄·5H₂O.

Water of crystallisation

The water molecules that are part of the crystal structure of a hydrated salt.

Anhydrous

A substance with no water of crystallisation, such as CuSO₄. It forms when the hydrated salt is heated.

Reacting masses and percentage yield

To find the mass of a product: (1) find the moles of the substance you know, (2) use the ratio from the balanced equation, (3) convert back to mass.

Yields are below 100% because the reaction may be reversible or incomplete, side reactions form other products, and some product is lost when it is separated and purified.

Volumes of gases

One mole of any gas occupies the same volume at the same temperature and pressure: 24.0 dm³ at room conditions, or 22.4 dm³ at s.t.p. So for gases, the ratio of volumes equals the ratio of moles in the equation.

Solutions and titrations

Concentration is the amount of solute in 1 dm³ of solution, in mol dm⁻³. Multiply by M to get g dm⁻³. Volumes measured in cm³ must be divided by 1000 to get dm³.

Limiting and excess reagents

The limiting reagent is the one that is completely used up. It decides how much product forms. The other reagent is in excess, and some of it is left over. To decide which is which, use the equation to find how much of one reactant is needed to react with all of the other.

Questions

Use A_r values: H 1.0, C 12.0, O 16.0, Mg 24.3, S 32.1, Ca 40.1. Questions 1 to 5 are multiple choice. Questions 6 to 8 are structured.

Bonds hold atoms together; intermolecular forces hold molecules together. This topic explains both, shows you how to draw dot-and-cross diagrams, and uses electron-pair repulsion to predict the shapes of molecules.

Electronegativity

Electronegativity

The power of an atom to attract electrons to itself (in a covalent bond).

Electronegativity depends on the same factors as ionisation energy: nuclear charge, atomic radius and shielding by inner shells and sub-shells. A small atom with a large nuclear charge and little shielding attracts bonding electrons strongly.

Fluorine is the most electronegative element. On the Pauling scale (values given in the exam when needed):

Predicting the type of bond

The difference in electronegativity between two atoms tells you how unevenly the bonding electrons are shared:

There is no sharp cut-off, so treat the boundary as a guide. Aluminium chloride is a useful case: its difference (1.55) is too small for ionic bonding, and it is indeed a covalent compound that forms Al₂Cl₆ molecules.

Ionic bonding

Ionic bonding

The electrostatic attraction between oppositely charged ions (positive cations and negative anions).

Ions form when a metal atom transfers its outer electrons to a non-metal atom, so that both end up with a noble gas configuration. In a dot-and-cross diagram, show only the outer shells, put each ion in square brackets, and write the charge at the top right.

In the solid, each ion is surrounded by ions of the opposite charge in a giant ionic lattice (Topic 4). Ionic bonding is stronger when the ions have larger charges and smaller radii, which is why magnesium oxide (2+ and 2−) melts at 2852 °C and sodium chloride (1+ and 1−) at 801 °C.

Metallic bonding

Metallic bonding

The electrostatic attraction between positive metal ions and delocalised electrons.

The outer electrons of the metal atoms are not tied to any one atom: they move freely through a lattice of positive ions (a "sea" of delocalised electrons). Metallic bonding is stronger when there are more delocalised electrons per atom and the ions are smaller and more highly charged. This explains why the melting points rise from Na (98 °C) to Mg (650 °C) to Al (660 °C), and why metals conduct electricity.

Covalent bonding

Covalent bond

The electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.

Each shared pair is one bond. Two shared pairs make a double bond and three make a triple bond. Pairs that are not shared are called lone pairs.

Expanding the octet

Elements in Period 2 can hold at most 8 electrons in their outer shell, because the second shell has only 2s and 2p orbitals. Elements from Period 3 onwards can use empty d orbitals, so they can have more than 8 outer electrons. Examples: sulfur dioxide (10 electrons around S), phosphorus(V) chloride PCl₅ (10 around P) and sulfur hexafluoride SF₆ (12 around S). This is why nitrogen forms only NCl₃ but phosphorus forms both PCl₃ and PCl₅.

Coordinate (dative covalent) bonding

Coordinate bond

A covalent bond in which both electrons in the shared pair come from the same atom.

The donor atom needs a lone pair, and the acceptor needs an empty orbital. Once formed, a coordinate bond is identical to any other covalent bond. It is often drawn as an arrow from the donor to the acceptor.

Ammonia and hydrogen chloride: the gases react to form white solid ammonium chloride. The lone pair on N forms a coordinate bond to H⁺: NH₃(g) + HCl(g) → NH₄Cl(s). All four N–H bonds in NH₄⁺ are then identical, and the ion is tetrahedral.

Aluminium chloride: in AlCl₃ the aluminium has only six outer electrons, so it has an empty orbital. In the vapour at lower temperatures, two AlCl₃ molecules join to form Al₂Cl₆: a lone pair on a chlorine atom of each molecule forms a coordinate bond to the aluminium of the other.

σ and π bonds

A covalent bond forms when orbitals on two atoms overlap, so that the shared pair is attracted to both nuclei.

• A σ (sigma) bond forms by direct, end-on overlap of orbitals between the two nuclei. Every single bond is a σ bond.
• A π (pi) bond forms by sideways overlap of adjacent p orbitals, with electron density above and below the σ bond. A double bond is one σ + one π; a triple bond is one σ + two π.

Hybridisation

Carbon's 2s and 2p orbitals mix (hybridise) to form a set of identical hybrid orbitals, which form the σ bonds. The number of hybrid orbitals equals the number of orbitals mixed.

In ethene, the π bond can only form if the p orbitals on the two carbons are parallel. So the molecule is planar and there is no rotation about C=C. This is the origin of cis/trans isomerism (Topic 13).

Bond energy and bond length

Bond energy

The energy required to break one mole of a particular covalent bond in the gaseous state.

Bond length

The internuclear distance between two covalently bonded atoms.

Some bond energies are exact: for a diatomic molecule such as H₂ there is only one bond in one environment. Others are averages: a C–H bond in methane has a slightly different energy from a C–H bond in ethanol, so the table gives a mean value.

• Multiple bonds are shorter and stronger: C≡C > C=C > C–C. But C=C is less than twice C–C, because the π bond is weaker than the σ bond.
• Longer bonds are generally weaker: the shared pair is further from the nuclei. H–Cl (0.127 nm, 431) is stronger than H–I (0.161 nm, 299), which is why hydrogen iodide decomposes much more easily (Topic 11).
• Strong bonds mean low reactivity: N≡N (944 kJ mol⁻¹) is one of the strongest bonds known, so nitrogen gas is very unreactive (Topic 12).

Shapes of molecules

Valence shell electron pair repulsion (VSEPR) theory: pairs of electrons around a central atom repel each other and move as far apart as possible. The shape depends on the number of bonding pairs and lone pairs.

• A double or triple bond counts as one region of electron density, just like a single bond.
• Lone pairs repel more strongly than bonding pairs, because they are closer to the central atom. The order is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. Each lone pair squeezes the bond angles by about 2.5°.
• The shape is named from the positions of the atoms, not the lone pairs.

More analogous examples: NH₂⁻ is non-linear (2 bonding, 2 lone pairs); CO₃²⁻ and NO₃⁻ are trigonal planar (3 regions, no lone pairs on the central atom); SO₂ is non-linear, about 120° (2 double bonds and 1 lone pair on S); PCl₄⁺ is tetrahedral; PCl₆⁻ is octahedral.

Intermolecular forces

Van der Waals' forces is the general name for all the attractive forces between molecules, other than those due to bond formation. There are two types in this course:

Instantaneous dipole–induced dipole (id-id) forces

Also called London dispersion forces, these act between all molecules, polar or not. Electrons are always moving, so at any instant the electron cloud may be unevenly spread, giving a temporary (instantaneous) dipole. This induces a dipole in a neighbouring molecule, and the two attract. The forces are stronger when:

• there are more electrons (larger molecules), so bigger dipoles form. Boiling points rise from F₂ (−188 °C) to Cl₂ (−34 °C), Br₂ (59 °C) and I₂ (184 °C).
• there is more surface contact between molecules. Pentane (straight chain, bp 36 °C) boils higher than its isomer 2,2-dimethylpropane (nearly spherical, bp 10 °C).

Permanent dipole–permanent dipole (pd-pd) forces

A polar molecule has a permanent dipole: one end δ+ and the other δ−. The δ+ end of one molecule attracts the δ− end of the next. These forces act in addition to id-id forces. For example, propanone (CH₃COCH₃, M_r 58, bp 56 °C) boils much higher than butane (C₄H₁₀, M_r 58, bp −1 °C), which has a similar number of electrons but no permanent dipole.

Hydrogen bonding

Hydrogen bonding is a special, strong case of permanent dipole–permanent dipole attraction. It happens between molecules where hydrogen is bonded to a highly electronegative atom. In this syllabus, that means molecules with N–H or O–H groups, such as ammonia, water and alcohols. (H–F also forms hydrogen bonds.)

1. The N–H or O–H bond is very polar, leaving the small H atom with a large δ+ charge.
2. This H is attracted to a lone pair on an N or O atom of a neighbouring molecule.
3. The three atoms O–H···O are in a straight line (180°), because this keeps the electron pairs as far apart as possible.

The unusual properties of water

Each water molecule has two H atoms and two lone pairs, so it can form up to four hydrogen bonds. This explains why water:

• has relatively high melting and boiling points. Hydrogen bonds are much stronger than the id-id forces in H₂S, so more energy is needed to separate the molecules. Water boils at 100 °C; H₂S, which has more electrons, boils at −60 °C.
• has a high surface tension. Molecules at the surface are pulled strongly inwards and sideways by hydrogen bonds.
• is less dense as ice than as liquid water. In ice, each molecule forms four hydrogen bonds in a regular tetrahedral arrangement, which holds the molecules further apart in an open lattice. When ice melts, some hydrogen bonds break and the molecules move closer together.

Bond polarity and polar molecules

A bond between atoms of different electronegativity is polar: the more electronegative atom gets a δ− charge. A molecule has an overall dipole moment (it is polar) only if its bond dipoles do not cancel out, which depends on its shape.

How strong are the forces?

In general, ionic, covalent and metallic bonds are much stronger than intermolecular forces. Breaking a covalent bond takes hundreds of kJ mol⁻¹. Hydrogen bonds are typically 5 to 40 kJ mol⁻¹, and other van der Waals' forces are usually weaker still. When a simple molecular substance melts or boils, only the intermolecular forces are overcome. The covalent bonds inside the molecules stay intact.

Dot-and-cross diagram checklist

• Show outer-shell electrons only, unless the question asks for all shells.
• Use dots for one atom's electrons and crosses for the other's, so the reader can see where each shared pair came from.
• Put shared pairs in the overlap between the two atoms, and include every lone pair.
• For ions, use square brackets and show the charge. For a coordinate bond, both electrons of the pair have the donor's symbol.
• Species with an odd number of electrons are allowed, such as the methyl radical CH₃•: C has 3 shared pairs and 1 unpaired electron.

Questions

Questions 1 to 6 are multiple choice. Questions 7 to 9 are structured.

The properties of a substance depend on its particles and on the forces between them. This topic covers how gases behave, how to use the ideal gas equation, and how the four types of structure explain melting points, conductivity and solubility.

The gaseous state

In a gas, the particles move rapidly and randomly in straight lines between collisions. Gas pressure comes from the particles colliding with the walls of the container: each collision exerts a force on the wall, and pressure is the total force per unit area.

Ideal gases

An ideal gas is a model that follows the gas laws exactly. Its key assumptions are:

• the particles have zero volume (they are points)
• there are no intermolecular forces of attraction between the particles
• (also) the particles move randomly and their collisions are perfectly elastic

Real gases

Real gases behave most like an ideal gas at high temperature and low pressure. They deviate when:

Gases with small, non-polar particles, such as helium and hydrogen, are closest to ideal. Gases with strong intermolecular forces, such as ammonia (hydrogen bonding) and steam, deviate the most.

The ideal gas equation

Using the ideal gas equation shows where the molar volume comes from: at 20 °C and 101 kPa, V = nRT/p = 1 × 8.31 × 293 / 101 000 = 0.0241 m³ = 24.1 dm³, close to the 24.0 dm³ used for room conditions.

Bonding and structure

Solids are classified by the particles in their lattice and the forces between them. There are four types:

Giant ionic: sodium chloride and magnesium oxide

A regular lattice of positive and negative ions held by strong ionic bonds in all directions. In sodium chloride, each Na⁺ ion is surrounded by six Cl⁻ ions and each Cl⁻ by six Na⁺ (6 : 6 coordination). Magnesium oxide has the same arrangement, with Mg²⁺ and O²⁻ ions.

Simple molecular: iodine, buckminsterfullerene and ice

Small molecules held in a lattice by weak intermolecular forces. The covalent bonds inside each molecule are strong, but only the weak forces between molecules are broken on melting.

• Iodine: I₂ molecules held by id-id forces. It forms shiny grey-black crystals that sublime on gentle warming to a purple vapour.
• Buckminsterfullerene, C₆₀: 60 carbon atoms joined in a hollow ball of 12 pentagons and 20 hexagons. The C₆₀ molecules are held by id-id forces, so it is soft, has a much lower sublimation point than diamond, does not conduct, and dissolves in non-polar solvents such as methylbenzene.
• Ice: H₂O molecules held in an open tetrahedral lattice by hydrogen bonds (Topic 3), which is why ice floats on water.

Giant molecular (giant covalent): diamond, graphite and silicon(IV) oxide

• Diamond: each carbon atom is covalently bonded to four others in a tetrahedral arrangement (109.5°), forming a 3D network. Very hard, very high melting point, does not conduct (no delocalised electrons).
• Graphite: each carbon is bonded to three others in flat layers of hexagons. The fourth outer electron of each carbon is delocalised over the layer, so graphite conducts electricity along the layers. The layers are held together by weak id-id forces, so they slide over each other: graphite is soft and slippery and is used as a lubricant and in pencils. Its melting point is still very high because the covalent bonds within the layers must break.
• Silicon(IV) oxide, SiO₂: each Si atom is bonded to four O atoms tetrahedrally, and each O atom to two Si atoms. Hard, high melting point (about 1700 °C), insoluble and non-conducting.

Giant metallic: copper

A regular lattice of positive copper ions in a sea of delocalised electrons. The metallic bonding is strong and non-directional, so copper has a high melting point (1085 °C). It conducts electricity because the delocalised electrons move through the structure. It is malleable and ductile because layers of ions can slide over each other without the metallic bonding breaking.

How structure explains properties

Questions

Use R = 8.31 J K⁻¹ mol⁻¹. Questions 1 to 4 are multiple choice. Questions 5 and 6 are structured.

Every reaction transfers energy to or from its surroundings, because bonds are broken and made. This topic covers how to define, measure and calculate enthalpy changes, including those you can't measure directly, using Hess's law.

Enthalpy changes

The enthalpy change, ΔH, is the heat energy transferred in a reaction at constant pressure, in kJ mol⁻¹.

Reaction pathway diagrams

A reaction pathway diagram (energy profile) shows the enthalpy of the reactants and products, and the energy barrier between them. The activation energy, E_A, is the minimum energy needed for a reaction to happen (Topic 8). It is measured from the reactants up to the top of the barrier.

Standard conditions and definitions

Enthalpy changes depend on conditions, so they are quoted under standard conditions, shown by the symbol : a temperature of 298 K and a pressure of 101 kPa, with solutions at 1 mol dm⁻³, and every substance in its standard state (its normal physical state under those conditions).

Enthalpy change of reaction, ΔH_r

The enthalpy change when the amounts of reactants shown in the equation react under standard conditions.

Enthalpy change of formation, ΔH_f

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.

Enthalpy change of combustion, ΔH_c

The enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions.

Enthalpy change of neutralisation, ΔH_neut

The enthalpy change when one mole of water is formed by the reaction of an acid with an alkali under standard conditions.

Bond energies

Breaking a bond needs energy (endothermic). Making a bond releases energy (exothermic). The overall enthalpy change depends on the balance between the two:

Measuring enthalpy changes

In calorimetry, the reaction heats (or cools) a known mass of water or solution, and you measure the temperature change.

m is the mass of the liquid being heated (for dilute solutions, take 1 cm³ = 1 g), c is 4.18 J g⁻¹ K⁻¹, ΔT is the temperature change, and n is the amount of the limiting reagent (or of water formed, for neutralisation). The minus sign makes ΔH negative when the temperature rises.

Correcting for heat loss

In Paper 3 you may record the temperature every minute before and after mixing. Extrapolating the cooling line back to the time of mixing gives the maximum temperature the mixture would have reached with no heat loss.

Hess's law

Hess's law

The total enthalpy change for a reaction is independent of the route taken, provided the starting and finishing conditions are the same.

Hess's law lets you calculate enthalpy changes that can't be measured directly, such as the enthalpy of formation of ethanol (you can't make ethanol by mixing carbon, hydrogen and oxygen). Draw a cycle with the reaction you want along the top and an alternative route through known enthalpy changes.

Using enthalpy changes of formation

Both reactants and products can be made from the same elements, so the arrows go up from the elements:

Using enthalpy changes of combustion

Both reactants and products burn to give the same combustion products, so the arrows go down to the combustion products:

Questions

Use c = 4.18 J g⁻¹ K⁻¹. Questions 1 to 4 are multiple choice. Questions 5 to 7 are structured.

Redox reactions transfer electrons. Oxidation numbers are the bookkeeping that tracks where the electrons go, and they let you name compounds, identify oxidising and reducing agents, and balance difficult equations.

Oxidation numbers

An oxidation number is the charge an atom would have if all its bonds were ionic. Work it out with these rules, in order:

Roman numerals in names give the oxidation number of the element before them: iron(III) chloride FeCl₃, potassium manganate(VII) KMnO₄, potassium dichromate(VI) K₂Cr₂O₇, sodium chlorate(I) NaClO, sodium chlorate(V) NaClO₃.

Redox

Remember "OIL RIG": Oxidation Is Loss, Reduction Is Gain (of electrons).

Disproportionation

Balancing redox equations

Method 1: half-equations

Balance each half separately, then add them so that the electrons cancel. In acidic solution, balance O atoms with H₂O and H atoms with H⁺, then balance charge with electrons.

Method 2: changes in oxidation number

The total increase in oxidation number must equal the total decrease.

Topic 24 (A Level) builds on this with electrode potentials, which predict whether a redox reaction will happen.

Questions

Many reactions never go to completion: they reach a balance in which reactants and products are both present. This topic explains that balance, how to shift it, how to calculate equilibrium constants, and how acids and bases behave.

Reversible reactions and dynamic equilibrium

A reversible reaction can go in both directions: the products can react to re-form the reactants. It is shown with the symbol ⇌, for example N₂(g) + 3H₂(g) ⇌ 2NH₃(g).

Dynamic equilibrium

The state in which the rate of the forward reaction equals the rate of the reverse reaction, so the concentrations of reactants and products stay constant.

"Dynamic" means both reactions are still happening; there is just no overall change. Equilibrium can only be reached in a closed system, where nothing can enter or leave. In an open system, a gaseous product would escape and the reverse reaction couldn't keep up.

Equilibrium is the same whichever side you start from: starting with only NH₃ gives the same equilibrium mixture as starting with the equivalent amounts of N₂ and H₂ under the same conditions.

Le Chatelier's principle

Le Chatelier's principle

If a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change.

Equilibrium constants

K_c, in terms of concentrations

For a general reaction aA + bB ⇌ cC + dD at equilibrium:

Square brackets mean concentration at equilibrium in mol dm⁻³. Products go on top, and each concentration is raised to the power of its coefficient. Work out the units from the expression: for N₂ + 3H₂ ⇌ 2NH₃, K_c = [NH₃]²[N₂][H₂]³ has units (mol dm⁻³)² / (mol dm⁻³)⁴ = mol⁻² dm⁶.

A large K_c means the equilibrium lies to the right (mostly products); a small one means it lies to the left.

K_p, in terms of partial pressures

Mole fraction

The amount of one gas divided by the total amount of gas: x_A = n_An_total

Partial pressure

The pressure that one gas in a mixture would exert if it occupied the container on its own: p_A = x_A × p_total. The partial pressures add up to the total pressure.

For gas-phase equilibria, K_p uses partial pressures in place of concentrations. For 2SO₂(g) + O₂(g) ⇌ 2SO₃(g): K_p = p_SO3²p_SO2² × p_O2, with units Pa⁻¹ (or kPa⁻¹).

In both kinds of calculation, set up a table of start, change and equilibrium amounts. At AS Level you won't need to solve a quadratic equation.

Equilibria in industry

Industrial conditions are a compromise between a good yield (position of equilibrium), a good rate, and cost and safety.

The Haber process

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹

The Contact process

2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = −197 kJ mol⁻¹. The SO₃ is then used to make sulfuric acid.

Acids and bases

Common acids and alkalis

The Brønsted–Lowry theory

Brønsted–Lowry acid

A proton (H⁺) donor.

Brønsted–Lowry base

A proton (H⁺) acceptor.

In HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq), HCl donates a proton (acid) and water accepts it (base). In NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq), water donates a proton (acid) and ammonia accepts it (base). Water can act as either, depending on what it reacts with.

Strong and weak acids and bases

"Strong" and "weak" describe the extent of dissociation, not the concentration. A dilute solution of a strong acid can have a lower [H⁺] than a concentrated solution of a weak one.

pH

The pH scale measures acidity: pure water has a pH of 7, acidic solutions are below 7 and alkaline solutions are above 7. Each unit of pH is a factor of 10 in the concentration of H⁺ ions (the pH calculations are in Topic 25).

Telling strong and weak acids apart

Compare solutions of the same concentration, for example 0.1 mol dm⁻³ HCl and 0.1 mol dm⁻³ CH₃COOH:

Neutralisation and salts

In every neutralisation between an acid and an alkali, H⁺(aq) + OH⁻(aq) → H₂O(l). The other ions form a salt, for example KOH(aq) + HNO₃(aq) → KNO₃(aq) + H₂O(l).

pH titration curves

A titration curve plots pH against the volume of alkali added. The shape depends on whether the acid and base are strong or weak. All four below use 25.0 cm³ of 0.100 mol dm⁻³ acid and 0.100 mol dm⁻³ base, so the equivalence point is at 25.0 cm³.

Choosing an indicator: its colour-change range must lie entirely within the steep (vertical) part of the curve. Then one drop of alkali at the end-point changes the pH right across the indicator's range.

Questions

Kinetics is about how fast reactions go. Collision theory and the Boltzmann distribution explain why concentration, pressure, temperature and catalysts change the rate.

Rate of reaction

Rate of reaction

The change in concentration of a reactant or product per unit time. Units: mol dm⁻³ s⁻¹.

Rates are measured by following a property that changes during the reaction: the volume of gas given off (gas syringe), the loss in mass as a gas escapes, a colour change (colorimeter), or by taking samples and titrating them.

• The average rate over a period = change in concentration ÷ time taken.
• The rate at a particular moment is the gradient of the tangent to the concentration–time curve at that time. The initial rate is the gradient at t = 0.

Collision theory

For particles to react, they must collide, with at least the activation energy and in the right orientation.

Activation energy, E_A

The minimum energy required for a collision to be effective.

Effective collision

A collision that results in a reaction.

Non-effective collision

A collision in which the particles bounce apart unchanged, because they have too little energy or the wrong orientation.

The rate depends on the frequency of effective collisions: the number per second.

Temperature and the Boltzmann distribution

In a gas (or solution), particles have a range of energies. The Boltzmann distribution shows how many particles have each energy.

Key features to show when you sketch the curve: it starts at the origin, has no particles with zero energy, rises to a peak (the most likely energy) and then falls towards the energy axis without touching or crossing it, because there is no maximum energy.

Why a higher temperature increases the rate:

1. Particles have more kinetic energy, so a much larger proportion of collisions have energy equal to or greater than E_A. This is the main effect.
2. Particles also move faster, so they collide more often. This is a much smaller effect.

Together, these increase the frequency of effective collisions. For many reactions near room temperature, a rise of 10 °C roughly doubles the rate.

Catalysts

Catalyst

A substance that increases the rate of a reaction without being used up overall. It provides an alternative reaction pathway (a different mechanism) with a lower activation energy.

With a lower E_A, a greater proportion of particles have enough energy to react, so the frequency of effective collisions increases. A catalyst does not change ΔH, and it does not change the position of equilibrium.

Homogeneous and heterogeneous catalysts

How these catalysts work in detail (adsorption, and catalysts that change oxidation state) is covered in Topic 26.

Questions

Period 3 runs from sodium to argon. Across it, the structure changes from giant metallic to giant covalent to simple molecular, and the oxides and chlorides change from ionic and basic to covalent and acidic. Learning why lets you predict the chemistry of any element from its position.

Physical properties across Period 3

Atomic and ionic radius

Atomic radius decreases from Na to Cl: the nuclear charge increases but the extra electrons go into the same (third) shell with similar shielding, so the outer electrons are pulled closer. For the ions, Na⁺ to Si⁴⁺ have lost their third shell and get smaller as nuclear charge rises; P³⁻ to Cl⁻ keep three shells, so they are much larger, and they also get smaller across the period. The values are in the table in Topic 1.

Melting point

Electrical conductivity

Na, Mg and Al are good conductors because of their delocalised electrons. Conductivity increases from Na to Al, as the number of delocalised electrons per atom increases. Silicon is a semiconductor: a poor conductor whose conductivity rises with temperature. Phosphorus, sulfur, chlorine and argon do not conduct, because all their electrons are held in covalent bonds or in single atoms.

Reactions of the elements

With oxygen

With chlorine

Sodium and magnesium with water

Oxidation numbers of the oxides and chlorides

The maximum oxidation number equals the number of outer-shell electrons, because in the highest oxide or chloride the element uses all of its outer electrons in bonding. In the chlorides it stops at +5 for phosphorus, because there isn't room around a sulfur atom for six chlorine atoms.

The oxides: acid–base behaviour

Reactions with acids and bases