Free A Level Chemistry 9701 Handouts

Topic 1

Atomic structure

Chemistry is the behaviour of electrons, so where the electrons sit decides everything an element does.

Picture itA stadium with seats arranged in rings. The seats nearest the pitch fill first and are hardest to leave. Every chemical property of an element — how readily it reacts, what it bonds to, what colour it is — comes down to which seats are occupied and how tightly the occupants are held.

Subatomic particles and isotopes

Protons and neutrons have relative mass 1; electrons about 1/1836. Charges are +1, 0 and −1. Isotopes have the same proton number but different neutron numbers, so they are chemically identical (same electron arrangement) but differ in mass — which is why mass spectrometry can separate them.

Subshells fill in a specific order

s holds 2 electrons, p holds 6, d holds 10, f holds 14. The filling order is 1s 2s 2p 3s 3p 4s 3d 4p — 4s fills before 3d because it is lower in energy when empty, but it is also emptied first when the atom ionises.

Ionisation energy and what it reveals

The first ionisation energy is the energy to remove one mole of electrons from one mole of gaseous atoms. Successive ionisation energies rise, with a large jump when a new shell is broken into — which is direct experimental evidence for shell structure and lets you deduce a group number.

The trends across and down

Across a period, ionisation energy rises as nuclear charge increases with similar shielding. It dips at group 13 (a p electron is easier to remove than an s electron) and at group 16 (paired p electrons repel). Down a group it falls, as the outer electron is further out and better shielded.

Orbital shapes

An s orbital is spherical; the three p orbitals are dumb-bell shaped along x, y and z. Each orbital holds at most two electrons, with opposite spins. Electrons occupy degenerate orbitals singly before pairing, which is why the group 16 dip exists at all.

The bit that catches people outThe 4s-before-3d order is not a rule about the atom in general — it is about relative energies at the point of filling. When a transition metal ionises, the 4s electrons leave first, which surprises students every single year.

The grown-up words

What it means	What it is called	Note
Atoms with the same protons, different neutrons	isotopes	Chemically identical
Energy to remove one mole of electrons from gaseous atoms	first ionisation energy	Endothermic
Region where an electron is likely to be found	orbital	Holds up to 2 electrons
Reduction of nuclear pull by inner electrons	shielding	Increases down a group
Orbitals of equal energy	degenerate orbitals	Fill singly before pairing
Charge felt by an outer electron	effective nuclear charge	Nuclear charge minus shielding

Check you have got it

Successive ionisation energies for an element jump sharply between the second and third. Which group is it in?

Group 2. Two electrons are removed relatively easily from the outer shell; the third must come from a full inner shell much closer to the nucleus.

Why is the first ionisation energy of aluminium lower than that of magnesium?

Aluminium's outer electron is in a 3p orbital, which is higher in energy and slightly better shielded than magnesium's 3s electron, so it is removed more easily.

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Topic 2

Atoms, molecules and stoichiometry

The mole exists because you cannot count atoms, but you can weigh them.

Picture itYou cannot count grains of rice for a recipe, so you weigh them. Chemists do the same: 12 g of carbon-12 contains 6.02 × 10²³ atoms, and every mole calculation you will ever do is a way of turning a mass you can measure into a number of particles you cannot.

The three core equations

n = m ÷ M for solids. n = cV for solutions (with V in dm³). n = V ÷ 24.0 for gases at room conditions, or pV = nRT in general with p in Pa, V in m³ and R = 8.31 J K⁻¹ mol⁻¹. Almost every stoichiometry question is one of these three, applied twice.

Empirical and molecular formulae

Divide each percentage or mass by the relative atomic mass, divide through by the smallest result, and scale to whole numbers — that is the empirical formula. Compare its mass with the relative molecular mass to find the multiple giving the molecular formula.

Limiting reagent and yield

Convert both reactants to moles, divide each by its coefficient, and the smaller value is limiting. Percentage yield = actual ÷ theoretical × 100. Atom economy = mass of desired product ÷ total mass of products × 100, which measures waste rather than efficiency of conversion.

Titration calculations

Use the mean of concordant titres — those within 0.10 cm³ of each other — and discard the rough. Then n = cV on the known solution, apply the mole ratio from the balanced equation, and solve for the unknown.

Balancing, including ionic and half equations

Balance atoms, then charges. For redox half equations: balance the main element, add H₂O for oxygen, H⁺ for hydrogen, then electrons to balance the charge. Combine the halves so the electrons cancel exactly.

The bit that catches people outConcentration is per dm³, but burettes and pipettes are marked in cm³. Dividing by 1000 at the right moment is the single most common arithmetic slip in the whole subject — write the units into the working every time.

The grown-up words

What it means	What it is called	Note
6.02 x 10^23 particles	one mole	Avogadro's constant
Simplest whole-number ratio of atoms	empirical formula	From composition data
Actual number of atoms in a molecule	molecular formula	A multiple of the empirical
Reactant fully used up first	limiting reagent	Determines the yield
Desired product mass over total product mass	atom economy	Measures waste
Titres within 0.10 cm3 of each other	concordant titres	Averaged for the calculation

Check you have got it

2.4 g of magnesium reacts with 100 cm³ of 1.0 mol dm⁻³ HCl. Which is limiting?

Mg: 2.4 ÷ 24.3 = 0.099 mol. HCl: 1.0 × 0.100 = 0.100 mol. The equation needs 2 HCl per Mg, so 0.099 mol Mg would require 0.198 mol HCl — the acid is limiting.

A compound is 40.0% C, 6.7% H, 53.3% O by mass. Find the empirical formula.

40.0/12.0 = 3.33; 6.7/1.0 = 6.7; 53.3/16.0 = 3.33. Dividing by 3.33 gives 1 : 2 : 1, so CH₂O.

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Topic 3

Chemical bonding

Atoms bond to reach a lower energy state, and the way they do it decides every physical property the substance has.

Picture itThree ways of holding a group together. Ionic is a transfer — one gives, one takes, and the opposite charges hold. Covalent is sharing — a pair of electrons held between two nuclei. Metallic is a communal pool — positive ions sitting in a sea of electrons that belong to nobody in particular.

The three bond types

Ionic: electrostatic attraction between oppositely charged ions in a giant lattice. Covalent: a shared pair attracted to both nuclei; a dative bond is one where both electrons come from the same atom. Metallic: cations in a delocalised electron sea, which explains conduction and malleability.

VSEPR predicts shape

Electron pairs repel and get as far apart as possible; lone pairs repel more strongly than bonding pairs. 2 pairs → linear 180°; 3 → trigonal planar 120°; 4 → tetrahedral 109.5°; 5 → trigonal bipyramidal; 6 → octahedral 90°. Each lone pair reduces the bond angle by roughly 2.5°.

Electronegativity and polarity

Unequal sharing gives a polar bond with δ+ and δ− ends. A molecule is only polar overall if the bond dipoles do not cancel — CO₂ is linear and symmetrical so it is non-polar, while H₂O is bent and is polar.

Intermolecular forces, weakest to strongest

Induced dipole (London) forces exist between all molecules and grow with the number of electrons. Permanent dipole forces act between polar molecules. Hydrogen bonding requires H attached to N, O or F and a lone pair on a neighbouring N, O or F.

Hydrogen bonding explains water's anomalies

Water's high boiling point, high surface tension, and — uniquely — ice being less dense than liquid water. In ice, each molecule hydrogen bonds to four others in an open tetrahedral arrangement that is more spacious than the liquid.

The bit that catches people outHydrogen bonding is not a bond between hydrogen atoms, and it does not happen just because a molecule contains hydrogen. It needs H bonded directly to N, O or F — which is why HCl does not hydrogen bond despite being polar.

The grown-up words

What it means	What it is called	Note
Electrostatic attraction between opposite ions	ionic bonding	Giant lattice
Shared pair of electrons	covalent bond	Attracted to both nuclei
Both electrons from the same atom	dative covalent bond	Also called coordinate
Cations in a sea of delocalised electrons	metallic bonding	Explains conduction
Model predicting molecular shape	VSEPR	Lone pairs repel most
Ability to attract a shared electron pair	electronegativity	Increases up and right
Force between all molecules from temporary dipoles	London forces	Grow with electron count
Strong dipole force with H bonded to N, O or F	hydrogen bonding	Explains water's anomalies

Check you have got it

Why is CO₂ non-polar despite having polar bonds?

The molecule is linear and symmetrical, so the two bond dipoles are equal and opposite and cancel exactly, leaving no overall dipole.

Predict the bond angle in NH₃ and explain it.

About 107°. There are four electron pairs around nitrogen, giving a tetrahedral arrangement, but one is a lone pair which repels more strongly than bonding pairs, squeezing the H–N–H angle below 109.5°.

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Topic 4

States of matter

Whether something is a gas, a liquid or a solid — and what kind of solid — comes down to what is holding the particles together.

Picture itMelting sugar and melting sand feel like the same kind of event, but they are not. Sugar melts at 186 °C because you only have to overcome forces between molecules. Sand melts above 1700 °C because you have to break the covalent bonds themselves. Same word, wildly different job.

The ideal gas equation and where it fails

pV = nRT assumes molecules have negligible volume and no intermolecular forces. Real gases deviate at high pressure (molecular volume is no longer negligible) and low temperature (attractions become significant). Ideal behaviour is best at low pressure and high temperature.

Simple molecular lattices

Iodine, ice, carbon dioxide: covalent bonds within molecules, weak intermolecular forces between them. Low melting points, poor conductors, often soluble in non-polar solvents. Melting breaks the intermolecular forces only — the molecules survive intact.

Giant covalent lattices

Diamond: each carbon bonded to four others tetrahedrally — very hard, non-conducting, extremely high melting point. Graphite: layers of hexagons with one delocalised electron per atom — conducts along layers, slippery because layers slide. Silicon dioxide resembles diamond.

Ionic and metallic lattices

Ionic lattices are hard, brittle and high melting; they conduct only when molten or dissolved, because the ions must be free to move. Metallic lattices are malleable because layers of cations can slide while the electron sea keeps holding them together.

Predicting properties from structure

Given an unknown, use melting point, conductivity as a solid and when molten, hardness and solubility to identify the structure. This deduction is a routine exam question and the reasoning must be explicit.

The bit that catches people outGraphite conducts and diamond does not, even though both are pure carbon. The reason is not the arrangement in the abstract — it is that in graphite each carbon uses only three of its four outer electrons for bonding, leaving one delocalised.

The grown-up words

What it means	What it is called	Note
Gas obeying pV = nRT exactly	ideal gas	Best at low p, high T
Molecules held by weak forces between them	simple molecular lattice	Low melting point
Network of covalent bonds throughout	giant covalent lattice	Very high melting point
Carbon bonded to four others tetrahedrally	diamond	Hard, non-conducting
Layers with one delocalised electron each	graphite	Conducts along layers
Ions in a regular repeating array	ionic lattice	Conducts when molten

Check you have got it

Why does an ionic solid conduct when molten but not when solid?

Conduction requires charged particles free to move. In the solid the ions are fixed in the lattice; melting frees them to move under a potential difference.

Under what conditions does a real gas most closely approach ideal behaviour, and why?

Low pressure and high temperature — molecules are far apart so their own volume is negligible, and they move fast enough that intermolecular attractions have little effect.

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Topic 5

Chemical energetics

Energy changes in reactions can be measured, and — more usefully — calculated from data you already have.

Picture itWalking between two towns. It does not matter whether you go directly or via three other villages: the change in altitude between start and finish is the same. That is Hess's law, and it is why you can calculate an enthalpy change for a reaction nobody has ever performed.

Enthalpy changes and their definitions

ΔH is negative for exothermic and positive for endothermic. Standard conditions: 298 K, 100 kPa, 1 mol dm⁻³. The definitions of enthalpy of formation, combustion, neutralisation and atomisation each specify exact amounts and states, and the marks are for that precision.

Hess's law and cycles

The enthalpy change is independent of route. Using enthalpies of formation: ΔH = Σ ΔH_f(products) − Σ ΔH_f(reactants). Using enthalpies of combustion: ΔH = Σ ΔH_c(reactants) − Σ ΔH_c(products). The two run in opposite directions, and mixing them up is the standard error.

Bond energies give approximate answers

ΔH ≈ Σ (bonds broken) − Σ (bonds made). Breaking is endothermic, making is exothermic. Values are only average bond energies across many compounds, so results are approximate — a point worth stating in the answer.

Calorimetry

q = mcΔT, where m is the mass of the solution being heated, not of the reactant. Divide by moles to get the enthalpy change per mole, and remember the sign: a temperature rise means an exothermic reaction and a negative ΔH.

Born–Haber cycles

Lattice energy cannot be measured directly, so it is found via a cycle combining atomisation, ionisation, electron affinity and enthalpy of formation. Lattice energy becomes more exothermic with higher ionic charge and smaller ionic radius.

The bit that catches people outHeat losses to the surroundings make experimental enthalpies of combustion less exothermic than the data book value, essentially every time. Incomplete combustion and evaporation of the fuel add to it. If a question asks why your value is 'too small', that is what it wants.

The grown-up words

What it means	What it is called	Note
Reaction releasing heat, negative dH	exothermic	Products lower in energy
Enthalpy change is route-independent	Hess's law	Basis of energy cycles
Energy to break one mole of a bond in the gas phase	bond energy	Average value
q = mcdT	calorimetry	m is the solution's mass
Energy released forming one mole of ionic solid from gaseous ions	lattice energy	Found via Born-Haber
Enthalpy change forming one mole from its elements	enthalpy of formation	Under standard conditions

Check you have got it

Calculate ΔH for a reaction where bonds broken total 2750 kJ and bonds made total 3100 kJ.

ΔH = 2750 − 3100 = −350 kJ mol⁻¹, so the reaction is exothermic.

Why is an experimental enthalpy of combustion usually less exothermic than the data book value?

Heat is lost to the surroundings and to the apparatus, combustion may be incomplete, and some fuel evaporates without burning — so less energy is transferred to the water than the reaction actually released.

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Topic 6

Electrochemistry (AS)

Redox is electron accounting, and oxidation numbers are the bookkeeping that makes it work.

Picture itA transaction ledger. Every electron that leaves one species arrives at another — none are created or destroyed. Oxidation numbers let you track them even in compounds where nothing obviously 'gave' anything, which is what makes balancing complicated redox equations possible.

Oxidation numbers and the rules

Elements are 0. Group 1 is +1, group 2 is +2. Oxygen is −2 (except peroxides −1 and OF₂ +2). Hydrogen is +1 (except metal hydrides −1). The sum equals the overall charge. Oxidation is an increase in oxidation number, reduction a decrease.

Half equations and combining them

Write each half equation with electrons, balance oxygen with H₂O and hydrogen with H⁺, then scale each half so the electrons cancel when added. This works for any redox equation, however awkward it looks.

Electrolysis

Cations migrate to the cathode and are reduced; anions migrate to the anode and are oxidised. Which species reacts depends on relative ease of discharge, concentration and the electrode material — which is why concentrated and dilute solutions of the same salt can give different products.

The Faraday relationships

Q = It, and the Faraday constant F = 96 500 C mol⁻¹ of electrons. Moles of electrons = It ÷ F; then apply the ratio from the half equation to get moles of product. Every quantitative electrolysis question is that chain.

Naming and predicting products

In aqueous solution, hydrogen is discharged in preference to reactive metals, and oxygen in preference to sulfate. Halides are discharged in preference to oxygen when concentrated — a case where concentration overrides the usual order.

The bit that catches people outOxidation is loss of electrons and an increase in oxidation number. Those two statements feel contradictory until you remember that losing negative charge makes the number go up. Most sign errors in this topic come from that one moment of confusion.

The grown-up words

What it means	What it is called	Note
Loss of electrons, oxidation number rises	oxidation	At the anode
Gain of electrons, oxidation number falls	reduction	At the cathode
Species causing another to be oxidised	oxidising agent	It is itself reduced
Equation showing electrons explicitly	half equation	Combined to cancel electrons
Decomposition using electricity	electrolysis	Cations to the cathode
96 500 coulombs per mole of electrons	Faraday constant	Q = It

Check you have got it

What is the oxidation number of manganese in MnO₄⁻?

+7. Four oxygens at −2 give −8; the overall charge is −1, so Mn must be +7.

A current of 2.0 A flows for 30 minutes through molten NaCl. How many moles of sodium form?

Q = 2.0 × 1800 = 3600 C. Moles of electrons = 3600 ÷ 96 500 = 0.0373. Na⁺ + e⁻ → Na is 1:1, so 0.037 mol of sodium.

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Topic 7

Equilibria (AS)

A reversible reaction settles at a balance point, and Le Chatelier's principle tells you which way that balance shifts when you disturb it.

Picture itA crowded room with two doors, people moving both ways. Once the rate through each door is equal, the number of people in each room stops changing — but nobody has stopped moving. That is dynamic equilibrium, and it is why 'the reaction stops' is always the wrong answer.

Dynamic equilibrium

Forward and reverse rates are equal, concentrations are constant, and the system must be closed. Nothing has stopped — the reaction continues in both directions at the same rate.

Le Chatelier's principle

A system at equilibrium opposes a change imposed on it. Increase concentration of a reactant and it shifts right. Increase pressure and it shifts to the side with fewer gas molecules. Increase temperature and it shifts in the endothermic direction. A catalyst changes neither position nor yield — only how fast equilibrium is reached.

The equilibrium constant

K_c = [products]^coefficients ÷ [reactants]^coefficients; K_p uses partial pressures. K changes only with temperature. Pressure and concentration changes shift the position but leave K untouched — a distinction worth several marks.

Industrial compromise

In the Haber process, low temperature favours yield but is too slow, so around 450 °C with an iron catalyst is a compromise. High pressure improves yield but costs money and safety, so about 200 atm is used. Recognising these as economic compromises, not chemical ideals, is the point.

Acids, bases and pH

Brønsted–Lowry: acids donate protons, bases accept them. Strong acids dissociate fully; weak acids partially. pH = −log[H⁺], and K_w = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 298 K. Conjugate acid–base pairs differ by one proton.

The bit that catches people outA catalyst does not increase yield. It speeds up the forward and reverse reactions equally, so equilibrium arrives sooner at exactly the same position. Writing that a catalyst 'increases the amount of product' is wrong every time.

The grown-up words

What it means	What it is called	Note
Forward and reverse rates equal	dynamic equilibrium	In a closed system
A system opposes an imposed change	Le Chatelier's principle	Predicts the shift
Ratio of product to reactant concentrations	Kc	Changes only with temperature
Equilibrium constant using partial pressures	Kp	For gaseous equilibria
Species differing by one proton	conjugate acid-base pair	Bronsted-Lowry
Ionic product of water	Kw	1.0 x 10^-14 at 298 K

Check you have got it

An exothermic equilibrium is heated. Which way does it shift and what happens to K?

It shifts in the endothermic direction — to the left — reducing the yield of products. K decreases, because K changes with temperature.

Why is the Haber process not run at a much lower temperature to maximise yield?

The forward reaction is exothermic, so a low temperature would give a higher equilibrium yield, but the rate would be far too slow to be economic. Around 450 °C is the compromise between yield and rate.

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Topic 8

Reaction kinetics (AS)

For a reaction to happen, particles must collide hard enough and in the right orientation — everything about rate follows from that.

Picture itThrowing two pieces of Velcro at each other. Most throws miss, some hit the wrong way round and bounce off, and only a few land hooks-to-loops with enough force to stick. Raising the temperature is throwing harder and more often; a catalyst is making the sticking easier.

Collision theory

A reaction requires a collision with energy at least equal to the activation energy, and with the correct orientation. Only a small fraction of collisions satisfy both, which is why reactions are far slower than collision frequency alone would suggest.

Factors affecting rate

Concentration and pressure increase collision frequency. Surface area increases the area available for collision. Temperature increases both frequency and — far more importantly — the proportion of collisions exceeding the activation energy. A catalyst provides an alternative route with a lower activation energy.

The Boltzmann distribution

A plot of number of molecules against energy, starting at the origin, peaking, and tailing off without touching the axis. Raising the temperature flattens and broadens the curve, shifting the peak right, so a much larger area lies beyond E_a — which is why a modest temperature rise has such a large effect.

Catalysts

Homogeneous catalysts are in the same phase as the reactants; heterogeneous catalysts are in a different phase and work by adsorbing reactants onto their surface. Neither is consumed overall, and neither changes ΔH or the equilibrium position.

Following a reaction experimentally

Measure gas volume, mass loss, colour change by colorimetry, or use titration on quenched samples. Rate is the gradient of the concentration–time graph; the initial rate is the gradient at t = 0.

The bit that catches people outA catalyst does not lower the activation energy of the original reaction — it provides an alternative pathway with a lower activation energy. The distinction sounds pedantic, but it is what the mark scheme asks for.

The grown-up words

What it means	What it is called	Note
Minimum energy needed for a successful collision	activation energy	Ea
Collision with enough energy and correct orientation	successful collision	Leads to reaction
Curve of molecule number against energy	Boltzmann distribution	Starts at the origin
Substance providing a lower-energy pathway	catalyst	Not consumed overall
Catalyst in the same phase as the reactants	homogeneous catalyst	e.g. aqueous ions
Catalyst in a different phase	heterogeneous catalyst	Works by adsorption

Check you have got it

Why does a 10 °C rise roughly double many reaction rates when it barely changes collision frequency?

Because the Boltzmann distribution shifts right, greatly increasing the proportion of molecules with energy above the activation energy. The increase in the number of successful collisions is far larger than the increase in total collisions.

Sketch how the Boltzmann curve changes when a catalyst is added.

The curve itself is unchanged — a catalyst does not alter molecular energies. Instead the activation energy line moves left, so a larger area under the same curve lies beyond it.

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Topic 9

Chemical periodicity

The periodic table is arranged so that properties repeat, and the repeats have causes you can state.

Picture itA spiral staircase. Every turn brings you back to the same position on the compass but one level higher. Sodium and potassium are the same 'position' — one outer electron — but on different levels, which is exactly why they behave alike and yet not identically.

Atomic radius and ionisation energy

Across period 3, atomic radius decreases: nuclear charge rises while shielding stays roughly constant, so the outer electrons are pulled in. Ionisation energy rises correspondingly, with the familiar dips at aluminium and sulfur.

Melting points across period 3

Na, Mg and Al are metallic and rise as charge and delocalised electron count increase. Si is giant covalent and is the highest of all. P, S and Cl are simple molecular, with S₈ higher than P₄ higher than Cl₂ because of London forces. Argon is monatomic and lowest.

Oxides across period 3

Na₂O and MgO are ionic and basic, giving alkaline solutions. Al₂O₃ is amphoteric, reacting with both acids and bases. SiO₂, P₄O₁₀, SO₂ and SO₃ are covalent and acidic, giving acidic solutions.

Chlorides across period 3

NaCl and MgCl₂ are ionic and dissolve to give near-neutral solutions. AlCl₃, SiCl₄, PCl₅ are covalent and hydrolyse vigorously in water, producing HCl fumes and strongly acidic solutions.

Group 2 and group 17 trends

Down group 2, reactivity increases as ionisation energy falls. Down group 17, reactivity decreases as the atoms get larger and less able to attract an electron — which is why chlorine displaces bromine from bromides.

The bit that catches people outSilicon dioxide is acidic even though it does not dissolve to give an acidic solution — it is classified by reacting with bases, not by its solubility. Insoluble does not mean unreactive, and the classification follows the chemistry.

The grown-up words

What it means	What it is called	Note
Repeating pattern of properties	periodicity	Down a group, across a period
Oxide reacting with both acids and bases	amphoteric oxide	e.g. Al2O3
Reaction with water splitting a compound	hydrolysis	Covalent chlorides fume
Distance from nucleus to outer electrons	atomic radius	Decreases across a period
Oxide giving an alkaline solution	basic oxide	Metals on the left
Oxide giving an acidic solution	acidic oxide	Non-metals on the right

Check you have got it

Why does silicon have the highest melting point in period 3?

It has a giant covalent structure, so melting requires breaking many strong covalent bonds throughout the lattice, rather than the weaker metallic or intermolecular forces in its neighbours.

Explain why AlCl₃ gives an acidic solution in water but NaCl does not.

AlCl₃ is covalent and hydrolyses, producing HCl and hydrated aluminium ions that release H⁺. NaCl is ionic and simply dissociates into hydrated ions that do not affect pH.

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Topic 10

Group 2 (AS)

Going down group 2, the elements get more reactive as metals and their compounds get harder to break down.

Picture itTwo trends running in the same direction for different reasons. The metals react more readily because their outer electrons are further out and easier to lose. Their carbonates become harder to decompose because the cations get bigger and distort the carbonate ion less.

Reactions with oxygen, water and acids

The metals burn in oxygen to give MO. They react with water — magnesium slowly with cold water but readily with steam, calcium onwards more vigorously — giving the hydroxide and hydrogen. With dilute acids they give a salt and hydrogen, more vigorously down the group.

Reactivity increases down the group

The outer electrons are further from the nucleus and better shielded, so ionisation energies fall and the metal loses electrons more readily. This is the same explanation as for group 1, and the marks are for stating both radius and shielding.

Hydroxide solubility increases down the group

Mg(OH)₂ is only sparingly soluble; Ba(OH)₂ is soluble. So the solutions become more alkaline down the group. Sulfate solubility runs the opposite way: MgSO₄ is soluble, BaSO₄ is essentially insoluble — which is why barium sulfate is the test for sulfate ions.

Thermal decomposition of carbonates and nitrates

Carbonates give the oxide and CO₂; nitrates give the oxide, NO₂ and O₂. Both become harder to decompose down the group, because larger cations have lower charge density and polarise the anion less, weakening it less.

Uses that follow from the chemistry

Ca(OH)₂ neutralises acidic soil. Mg(OH)₂ and CaCO₃ treat indigestion. BaSO₄ is a radiocontrast agent — safe only because it is so insoluble that essentially no toxic barium ions enter the body.

The bit that catches people outSolubility of hydroxides increases down the group while solubility of sulfates decreases. Two opposite trends, easily confused under exam pressure — anchor them with barium: Ba(OH)₂ dissolves, BaSO₄ does not.

The grown-up words

What it means	What it is called	Note
Charge divided by ionic radius	charge density	Falls down the group
Distortion of an anion's electron cloud	polarisation	Greater with small cations
Breaking down on heating	thermal decomposition	Harder down group 2
Sparingly soluble group 2 hydroxide	magnesium hydroxide	Used in antacids
Insoluble group 2 sulfate	barium sulfate	Test for sulfate ions
Reducing soil acidity with a base	liming	Uses calcium hydroxide

Check you have got it

Why is barium carbonate harder to decompose than magnesium carbonate?

The Ba²⁺ ion is much larger, so it has a lower charge density and polarises the carbonate ion less. The C–O bonds are weakened less, so more energy is needed to break the carbonate down.

Why is barium sulfate safe to swallow when barium ions are toxic?

Barium sulfate is essentially insoluble, so almost no free Ba²⁺ ions are released into the body.

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Topic 11

Group 17

The halogens all want one more electron, and how badly they want it decreases neatly down the group.

Picture itA queue for one remaining seat. Fluorine is closest to the front and takes it instantly. Iodine is at the back and often misses out. That single ordering explains displacement reactions, oxidising power, and why the hydrogen halides get less thermally stable downwards.

Physical trends

Down the group, molecules get larger with more electrons, so London forces strengthen: chlorine is a gas, bromine a liquid, iodine a solid. Colour deepens correspondingly.

Oxidising power decreases down the group

Larger atoms with more shielding attract an extra electron less strongly. So chlorine displaces bromide and iodide; bromine displaces only iodide; iodine displaces neither. These displacement reactions are the standard practical demonstration.

Reactions with hydrogen and thermal stability

HF, HCl, HBr, HI form less readily and are less thermally stable down the group, because the bond gets longer and weaker as the halogen atom gets bigger. HI decomposes noticeably on gentle heating.

Disproportionation of chlorine

With cold dilute NaOH, chlorine gives chloride and chlorate(I) — used in bleach. With hot concentrated NaOH it gives chloride and chlorate(V). In both, chlorine is simultaneously oxidised and reduced, which is what disproportionation means.

Testing for halide ions

Add dilute nitric acid then silver nitrate: chloride gives a white precipitate that dissolves in dilute ammonia; bromide gives cream, dissolving only in concentrated ammonia; iodide gives yellow, insoluble in concentrated ammonia. The ammonia step is what makes the test conclusive.

The bit that catches people outAdding chlorine to drinking water is a genuine risk-benefit judgement, not a straightforward good. It kills pathogens that have killed enormous numbers of people, and it can form small quantities of chlorinated organic compounds. The exam expects you to weigh both, not to pick a side.

The grown-up words

What it means	What it is called	Note
Same element oxidised and reduced	disproportionation	Chlorine with NaOH
Ability to accept electrons	oxidising power	Decreases down group 17
More reactive halogen replacing a less reactive one	displacement reaction	Cl2 displaces Br-
White precipitate with silver nitrate	chloride ion	Dissolves in dilute ammonia
Cream precipitate with silver nitrate	bromide ion	Needs concentrated ammonia
Yellow precipitate with silver nitrate	iodide ion	Insoluble in concentrated ammonia

Check you have got it

Why does the thermal stability of the hydrogen halides decrease from HCl to HI?

The halogen atom gets larger, so the H–X bond is longer and weaker, requiring less energy to break.

Write what happens when chlorine reacts with cold dilute sodium hydroxide.

Cl₂ + 2NaOH → NaCl + NaClO + H₂O. Chlorine is disproportionated from 0 to −1 in chloride and +1 in chlorate(I).

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Topic 12

Nitrogen and sulfur

Two elements whose industrial chemistry feeds the world and whose escaped emissions damage it.

Picture itNitrogen makes up 78% of the air and almost nothing can use it, because the triple bond is extraordinarily strong. The Haber process breaks that bond and turns air into fertiliser — arguably the most consequential chemical process ever industrialised, and one with genuine environmental costs.

Why nitrogen is unreactive

The N≡N triple bond has a bond energy around 945 kJ mol⁻¹ and the molecule is non-polar. Breaking it requires extreme conditions, which is why the Haber process needs high pressure, moderately high temperature and an iron catalyst.

Ammonia and ammonium salts

Ammonia is a base: its lone pair accepts a proton to form NH₄⁺ by a dative bond. It is a weak base in water. Warming an ammonium salt with alkali releases ammonia gas, which turns damp red litmus blue — the standard test.

Oxides of nitrogen

NO and NO₂ form in car engines at high temperature. NO₂ contributes to acid rain and to photochemical smog, and nitrogen oxides catalyse the oxidation of SO₂ to SO₃ in the atmosphere. Catalytic converters reduce them back to nitrogen.

Sulfur and the contact process

SO₂ is oxidised to SO₃ over vanadium(V) oxide at about 450 °C and 1–2 atm, then absorbed into concentrated sulfuric acid and diluted. Directly dissolving SO₃ in water is avoided because the reaction is dangerously vigorous.

Acid rain and its control

SO₂ from fossil fuels dissolves to give sulfurous and sulfuric acid, damaging buildings, forests and freshwater ecosystems. Flue gas desulfurisation with calcium oxide or calcium carbonate removes SO₂ before it is released, producing calcium sulfate.

The bit that catches people outFertiliser is not simply a good thing or a bad thing. Nitrogenous fertiliser massively raised food production; run-off into rivers causes eutrophication, where algal blooms deoxygenate water and kill fish. A strong answer holds both facts at once.

The grown-up words

What it means	What it is called	Note
Very strong bond making N2 unreactive	triple bond	About 945 kJ per mole
Base accepting a proton via its lone pair	ammonia	Forms NH4+ datively
Industrial synthesis of ammonia	Haber process	Iron catalyst, about 450 C
Industrial synthesis of sulfuric acid	contact process	Vanadium(V) oxide catalyst
Removal of SO2 from power station emissions	flue gas desulfurisation	Uses calcium compounds
Nutrient enrichment killing aquatic life	eutrophication	From fertiliser run-off

Check you have got it

Why is SO₃ not added directly to water in the contact process?

The reaction is violently exothermic and produces a dense acid mist that is difficult to condense and hazardous. Absorbing SO₃ into concentrated sulfuric acid first, then diluting, controls it.

Explain how fertiliser run-off can kill fish.

Nitrates cause rapid algal growth; when the algae die, aerobic bacteria decomposing them consume dissolved oxygen, leaving too little for fish — eutrophication.

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Topic 13

An introduction to AS organic chemistry

Organic chemistry is not thousands of unrelated reactions — it is a small number of functional groups behaving predictably.

Picture itA language with a modest vocabulary and strict grammar. Once you know what an alkene does, you know what every alkene does. The molecule around the functional group mostly just changes the name, not the behaviour.

Naming: the rules are mechanical

Find the longest carbon chain for the stem; identify the principal functional group for the suffix; number so the group gets the lowest locant; list substituents alphabetically with their numbers. Applied strictly, naming has one right answer.

Formulae and representations

Empirical, molecular, structural, displayed and skeletal formulae each show a different amount of detail. Skeletal formulae are fastest for large molecules and are expected at A Level.

Structural isomerism

Chain isomers differ in the carbon skeleton, positional isomers in where the group sits, functional group isomers in which group is present entirely — as with an aldehyde and a ketone of the same molecular formula.

Stereoisomerism

Cis–trans (E/Z) isomerism arises because the C=C double bond cannot rotate, so groups are locked on the same or opposite sides. Optical isomerism arises from a chiral carbon with four different groups, giving non-superimposable mirror images that rotate plane-polarised light in opposite directions.

Bond breaking and reagent types

Homolytic fission gives two radicals; heterolytic fission gives an anion and a cation. Electrophiles accept an electron pair; nucleophiles donate one. Curly arrows always show the movement of an electron pair, starting from a bond or a lone pair.

The bit that catches people outA curly arrow shows where an electron pair moves from and to. Starting an arrow at a positive charge, or at an atom rather than at a bond or lone pair, is wrong even when the overall answer is right — and mechanism marks are given for exactly this.

The grown-up words

What it means	What it is called	Note
Atom group giving a molecule its characteristic reactions	functional group	Determines behaviour
Series differing by CH2 with the same general formula	homologous series	Gradual property change
Same molecular formula, different arrangement	structural isomerism	Chain, positional, functional
Isomerism from restricted rotation about C=C	cis-trans isomerism	Also written E/Z
Carbon with four different groups attached	chiral centre	Gives optical isomers
Bond breaking giving two radicals	homolytic fission	One electron each
Bond breaking giving two ions	heterolytic fission	Both electrons to one atom
Electron pair acceptor	electrophile	Nucleophile is the donor

Check you have got it

Why can but-2-ene show cis–trans isomerism but but-1-ene cannot?

Cis–trans isomerism needs two different groups on each carbon of the C=C. In but-1-ene one double-bonded carbon carries two hydrogens, so swapping them changes nothing.

Identify the functional group isomer of propanal and name it.

Propanone — both are C₃H₆O, but propanal is an aldehyde and propanone is a ketone.

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Topic 14

Hydrocarbons (AS)

Alkanes are unreactive and burn; alkenes have a reactive double bond that everything wants to attack.

Picture itA closed fist and an open hand. An alkane's electrons are all tied up in strong, non-polar single bonds — nothing to grab. An alkene's π bond sticks out above and below the molecule, an accessible region of high electron density that electrophiles head straight for.

Alkanes: combustion and substitution

Complete combustion gives CO₂ and water; incomplete combustion gives CO and soot. With halogens in UV light, alkanes undergo free radical substitution: initiation (homolytic fission of the halogen), propagation (two steps that regenerate a radical), and termination (two radicals combine).

Why alkanes are otherwise unreactive

C–C and C–H bonds are strong and essentially non-polar, so there is no δ+ site for a nucleophile and no accessible electron density for an electrophile.

Alkenes: electrophilic addition

The π bond attacks an electrophile, forming a carbocation intermediate, which is then attacked by a nucleophile. Reagents: H₂/Ni, Br₂ (the decolourising test), HBr, H₂O/H₃PO₄ or steam, and acidified KMnO₄ giving a diol.

Markovnikov's rule and carbocation stability

When HBr adds to an unsymmetrical alkene, the major product comes from the more stable carbocation. Tertiary is more stable than secondary, which is more stable than primary, because alkyl groups are electron-releasing and spread the positive charge.

Addition polymerisation

Alkene monomers open their double bonds and link into long chains — poly(ethene), poly(propene), PVC. The saturated backbone is chemically inert, which is useful and is also exactly why disposal is a problem: landfill persistence, and toxic products such as HCl if PVC is burned carelessly.

The bit that catches people outBromine water decolourises with an alkene but not with an alkane. That is the standard test — but note that it works because of an addition reaction, whereas an alkane would only react with bromine under UV light, and by substitution.

The grown-up words

What it means	What it is called	Note
Reaction where a radical replaces an atom	free radical substitution	Alkanes with UV light
Species with an unpaired electron	free radical	Formed by homolytic fission
Adding across a double bond	addition reaction	Characteristic of alkenes
Positively charged carbon intermediate	carbocation	Tertiary most stable
Predicts the major product of unsymmetrical addition	Markovnikov's rule	Via the stabler carbocation
Monomers joining with no small molecule lost	addition polymerisation	e.g. poly(ethene)

Check you have got it

HBr adds to propene. Which is the major product and why?

2-bromopropane. Adding H⁺ to the terminal carbon gives a secondary carbocation, which is more stable than the primary alternative, so that route dominates.

Give the three stages of free radical substitution and one problem with the mechanism.

Initiation, propagation, termination. A problem is that it gives a mixture — further substitution produces di- and tri-halogenated products, so the yield of any single product is poor.

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Topic 15

Halogen compounds (AS)

A halogen on a carbon makes that carbon slightly positive — and a slightly positive carbon is an invitation to nucleophiles.

Picture itA magnet with a weak positive end. The halogen is more electronegative, pulling electron density away and leaving the carbon δ+. Anything with a lone pair heads for it. That single fact generates the whole reaction set for this topic.

Nucleophilic substitution

With aqueous NaOH → alcohol. With ethanolic KCN → nitrile, gaining a carbon. With ethanolic NH₃ under pressure → amine. Each is a nucleophile attacking the δ+ carbon and displacing the halide.

SN1 versus SN2

SN2 is a one-step attack from behind, with inversion of configuration — favoured by primary halogenoalkanes. SN1 proceeds via a carbocation intermediate — favoured by tertiary, because the carbocation is stabilised. Secondary can go either way.

Elimination competes with substitution

Hot ethanolic hydroxide favours elimination, giving an alkene. Warm aqueous hydroxide favours substitution, giving an alcohol. Same reagent, different solvent, different product — a classic exam distinction.

Rate depends on the C–X bond strength, not polarity

C–I is the weakest bond and reacts fastest; C–F is the strongest and slowest. This runs opposite to the polarity trend, which would predict fluorine fastest — bond enthalpy wins, and saying so explicitly earns the mark.

Uses and environmental effects

Halogenoalkanes are useful as solvents, refrigerants and anaesthetics. CFCs release chlorine radicals in the stratosphere which catalytically destroy ozone — one radical destroys many ozone molecules before termination, which is why small quantities mattered so much.

The bit that catches people outThe reactivity order for hydrolysis is iodo > bromo > chloro, which is the opposite of what bond polarity suggests. The C–I bond is weakest and breaks most easily, and bond strength dominates. Arguing from polarity alone gets the answer backwards.

The grown-up words

What it means	What it is called	Note
Electron pair donor attacking a positive centre	nucleophile	OH-, CN-, NH3
One-step substitution with inversion	SN2	Favoured by primary
Substitution via a carbocation	SN1	Favoured by tertiary
Removal of HX to form a double bond	elimination	Hot ethanolic hydroxide
Reaction with water breaking a bond	hydrolysis	Aqueous hydroxide
Chlorine-containing compound damaging ozone	CFC	Radicals act catalytically

Check you have got it

Why does hot ethanolic KOH give a different product from warm aqueous KOH?

In ethanol the hydroxide acts as a base, removing a hydrogen from the adjacent carbon and eliminating HX to give an alkene. In water it acts as a nucleophile, substituting to give an alcohol.

Explain why one chlorine radical can destroy many ozone molecules.

It acts catalytically: it reacts with ozone to form ClO and O₂, then ClO reacts with an oxygen atom to regenerate the chlorine radical, which attacks another ozone molecule.

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Topic 16

Hydroxy compounds (AS)

The OH group hydrogen bonds, which makes alcohols behave far less like their parent alkanes than their size suggests.

Picture itEthane boils at −89 °C. Ethanol, almost the same size and mass, boils at 78 °C. The entire difference is hydrogen bonding between the OH groups. One functional group, 167 degrees of consequence.

Physical properties

Alcohols have much higher boiling points than comparable alkanes and are soluble in water for short chains, both because of hydrogen bonding. Solubility falls as the non-polar chain lengthens and dominates.

Classification matters for oxidation

Primary alcohols oxidise to an aldehyde (distil off immediately) and then to a carboxylic acid (reflux). Secondary alcohols oxidise to a ketone. Tertiary alcohols resist oxidation, because there is no hydrogen on the carbon bearing the OH.

The other reactions

With sodium → alkoxide and hydrogen. With carboxylic acids and an acid catalyst → ester (esterification). With concentrated H₂SO₄ or Al₂O₃, hot → dehydration to an alkene. With PCl₅, PCl₃ or SOCl₂ → chloroalkane.

The tri-iodomethane test

Iodine and alkali give a pale yellow precipitate of CHI₃ with methyl ketones and with alcohols containing the CH₃CH(OH)– group. It is a positive test for a specific structural fragment, not for alcohols in general.

Making ethanol two ways

Hydration of ethene: fast, pure product, continuous, but uses a finite fossil feedstock. Fermentation: renewable and low-technology, but slow, batch, and gives a dilute product needing distillation.

The bit that catches people outDistillation and reflux are not interchangeable. To stop at the aldehyde you distil it out as it forms, before it can be oxidised further. To reach the carboxylic acid you reflux, deliberately returning everything to the flask. The apparatus is the chemistry.

The grown-up words

What it means	What it is called	Note
Alcohol with OH on a carbon bearing two hydrogens	primary alcohol	Oxidises to aldehyde then acid
Alcohol with OH on a carbon bearing one hydrogen	secondary alcohol	Oxidises to a ketone
Alcohol resistant to oxidation	tertiary alcohol	No H on the OH carbon
Heating with continuous condensation and return	reflux	Drives oxidation to the acid
Removal of water to form an alkene	dehydration	Concentrated H2SO4
Alcohol plus acid giving a sweet-smelling product	esterification	Acid catalysed
Pale yellow precipitate with iodine and alkali	tri-iodomethane test	Detects CH3CH(OH) group

Check you have got it

Why can a tertiary alcohol not be oxidised by acidified dichromate?

Oxidation requires removing a hydrogen from the carbon bearing the OH group. In a tertiary alcohol that carbon has no hydrogen, so oxidation would require breaking a C–C bond.

Give one advantage each of fermentation and of hydration of ethene for making ethanol.

Fermentation uses a renewable feedstock and needs only simple, low-energy equipment. Hydration of ethene is fast, continuous and gives a pure product needing no purification.

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Topic 17

Carbonyl compounds

The C=O bond is polar, so unlike the alkene's C=C it attracts nucleophiles rather than electrophiles.

Picture itTwo double bonds that look alike on paper and behave in opposite ways. In C=C the electrons are shared evenly, so electrophiles attack. In C=O oxygen pulls the electrons over, leaving the carbon δ+ — so nucleophiles attack instead. One difference in electronegativity, two entirely different chemistries.

Aldehydes and ketones

An aldehyde has the carbonyl at the end of the chain, with a hydrogen attached; a ketone has it in the middle. That hydrogen is why aldehydes are readily oxidised and ketones are not.

Nucleophilic addition

With HCN (with a trace of KCN), the cyanide ion attacks the δ+ carbon to give a hydroxynitrile, extending the carbon chain by one. With NaBH₄ or LiAlH₄, hydride reduces aldehydes to primary alcohols and ketones to secondary alcohols.

Distinguishing aldehydes from ketones

Tollens' reagent: aldehydes give a silver mirror. Fehling's solution: aldehydes give a brick-red precipitate of Cu₂O. Ketones do neither. Both are oxidation tests, and only aldehydes are oxidisable.

Detecting a carbonyl group

2,4-dinitrophenylhydrazine gives an orange or yellow crystalline precipitate with any aldehyde or ketone. Recrystallising it and measuring its melting point identifies the specific carbonyl compound — an elegant piece of classical analysis.

The tri-iodomethane test again

Methyl ketones — those with a CH₃C=O group — give the pale yellow CHI₃ precipitate with iodine and alkali, which distinguishes them from other ketones.

The bit that catches people outNucleophiles attack C=O; electrophiles attack C=C. Applying alkene reasoning to a carbonyl, or the reverse, produces a confident and completely wrong mechanism — check which double bond you are looking at before you draw a single arrow.

The grown-up words

What it means	What it is called	Note
Carbonyl at the end of a chain	aldehyde	Readily oxidised
Carbonyl in the middle of a chain	ketone	Resists oxidation
Nucleophile adds across the C=O bond	nucleophilic addition	With HCN or hydride
Silver mirror with an aldehyde	Tollens' reagent	Ketones give no reaction
Brick-red precipitate with an aldehyde	Fehling's solution	Copper(I) oxide
Orange precipitate with any carbonyl	2,4-DNPH	Identified by melting point

Check you have got it

Explain why aldehydes give a silver mirror with Tollens' reagent but ketones do not.

Aldehydes have a hydrogen on the carbonyl carbon and can be oxidised to a carboxylic acid, reducing the silver ions to silver. Ketones would require breaking a C–C bond, so they are not oxidised.

What is the product of reacting propanone with HCN, and why is the carbon chain longer?

2-hydroxy-2-methylpropanenitrile. The cyanide ion attacks the carbonyl carbon, adding a new carbon atom as part of the nitrile group.

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Topic 18

Carboxylic acids and derivatives (AS)

The COOH group is a weak acid, and the various things you can replace its OH with are its derivatives.

Picture itA hub with spokes. Carboxylic acid at the centre; replace the OH with OR and you have an ester, with Cl an acyl chloride, with NH₂ an amide. Every derivative can be made from the acid and hydrolysed back to it, so learning the hub gives you the whole wheel.

Acidity and why it exists

Carboxylic acids are weak — only partially dissociated. The carboxylate ion is stabilised by delocalisation of the negative charge over both oxygens, which is why they are far more acidic than alcohols. Electron-withdrawing groups nearby increase acidity further.

Reactions confirming acidity

With carbonates → CO₂ (effervescence, a distinguishing test from phenols and alcohols). With reactive metals → salt and hydrogen. With bases → salt and water. With alcohols and an acid catalyst → esters.

Esters

Made by esterification (reversible, acid catalysed) or from an acyl chloride (fast, essentially irreversible). Hydrolysed by acid (reversible, gives the acid and the alcohol) or by alkali (irreversible, gives the carboxylate salt and the alcohol). Used as flavourings, solvents and plasticisers.

Acyl chlorides

Made from the acid with PCl₅, PCl₃ or SOCl₂. Extremely reactive: they react vigorously with water, alcohols, ammonia and amines, giving the acid, ester, amide and substituted amide, releasing HCl each time.

Triglycerides and saponification

Fats are esters of glycerol with three fatty acids. Alkaline hydrolysis — saponification — gives glycerol and the sodium salts of the fatty acids, which are soaps. Biodiesel is made by transesterification with methanol.

The bit that catches people outAcid hydrolysis of an ester is reversible and never goes to completion; alkaline hydrolysis is irreversible because the carboxylate ion formed will not react back with the alcohol. That is why soap-making uses alkali.

The grown-up words

What it means	What it is called	Note
Partially dissociated acid	weak acid	Carboxylic acids
Spreading of charge over several atoms	delocalisation	Stabilises the carboxylate ion
Acid plus alcohol with an acid catalyst	esterification	Reversible
Highly reactive COCl compound	acyl chloride	Releases HCl
Alkaline hydrolysis of a fat	saponification	Gives soap and glycerol
Ester of glycerol and three fatty acids	triglyceride	The structure of fats

Check you have got it

Why is ethanoic acid a stronger acid than ethanol?

The ethanoate ion formed on dissociation is stabilised by delocalisation of the negative charge across both oxygen atoms, so the equilibrium lies further towards dissociation. The ethoxide ion has no such stabilisation.

How would you distinguish a carboxylic acid from a phenol using a simple test?

Add sodium carbonate solution. The carboxylic acid effervesces, releasing CO₂; phenol is too weak an acid to react.

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Topic 19

Nitrogen compounds (AS)

Amines are bases because nitrogen keeps a lone pair, and how available that lone pair is decides how strong a base you have.

Picture itA hand held out to shake. Nitrogen's lone pair is the hand — available to take a proton. Attach electron-releasing alkyl groups and the hand is pushed further forward, so the base is stronger. Attach a benzene ring that pulls the pair into itself, and the hand is withdrawn.

Amines as bases

The lone pair accepts a proton to give a substituted ammonium ion. Alkyl groups are electron-releasing, so aliphatic amines are stronger bases than ammonia. In phenylamine the lone pair is delocalised into the ring, making it a much weaker base.

Making amines

From a halogenoalkane with excess ethanolic ammonia under pressure — but the product is itself nucleophilic, so further substitution gives a mixture. From a nitrile by reduction with LiAlH₄ or H₂/Ni. Phenylamine is made by reducing nitrobenzene with tin and concentrated HCl.

Amino acids

Contain both COOH and NH₂, so they are amphoteric. In solution they exist as zwitterions, internally charged, which explains their high melting points and their solubility in water. The isoelectric point is the pH at which the zwitterion has no net charge.

Peptides and proteins

The COOH of one amino acid condenses with the NH₂ of another, forming a peptide bond and releasing water. Acid or alkaline hydrolysis breaks proteins back into their constituent amino acids.

Chromatography and analysis

Amino acid mixtures are separated by paper or thin-layer chromatography, identified by comparing R_f values with known standards, and visualised with ninhydrin because they are colourless.

The bit that catches people outPhenylamine is a much weaker base than ethylamine even though both are amines. The lone pair on nitrogen is delocalised into the benzene ring and so is far less available to accept a proton — structure decides the property.

The grown-up words

What it means	What it is called	Note
Molecule with both positive and negative charges	zwitterion	Amino acids in solution
Reacting as both an acid and a base	amphoteric	Because of COOH and NH2
pH at which there is no net charge	isoelectric point	Amino acid does not migrate
Bond joining two amino acids	peptide bond	Formed by condensation
Distance moved by solute over solvent	Rf value	Identifies a compound
Reagent revealing colourless amino acids	ninhydrin	Gives purple spots

Check you have got it

Why do amino acids have unexpectedly high melting points for their size?

They exist as zwitterions with full positive and negative charges, so there are strong ionic attractions between molecules in the solid, requiring much more energy to overcome than ordinary intermolecular forces.

Why does reacting a halogenoalkane with ammonia give a mixture of products?

The primary amine formed is itself a nucleophile and attacks further halogenoalkane, producing secondary and tertiary amines and eventually the quaternary ammonium salt.

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Topic 20

Polymerisation (AS)

Long molecules made by joining small ones — and how they were joined decides whether they can ever be unjoined.

Picture itTwo ways of making a chain. Addition polymers are made by opening double bonds — nothing is lost, and the resulting backbone is all C–C, which nothing in nature can break. Condensation polymers lose a small molecule at each join, and that join can be hydrolysed back.

Addition polymers

Alkene monomers open their C=C and link. Poly(ethene), poly(propene), poly(chloroethene) (PVC), PTFE. The repeat unit is drawn with the double bond opened and bonds extending through the brackets. The saturated backbone makes them chemically inert.

Condensation polymers

Polyesters (Terylene) from a diol and a dicarboxylic acid, losing water. Polyamides (nylon, Kevlar) from a diamine and a dicarboxylic acid or diacyl chloride. Proteins are natural polyamides.

Drawing repeat units and identifying monomers

For a condensation polymer shown as a chain, cut at the ester or amide link and add back OH and H to recover the monomers. This is a standard question and it rewards careful drawing rather than memory.

Properties from structure

Kevlar's strength comes from extensive hydrogen bonding between aligned chains. Chain length, branching and cross-linking determine strength and melting behaviour. Plasticisers work by pushing chains apart, weakening intermolecular forces.

Disposal is a chemistry problem

Addition polymers are non-biodegradable and persist for centuries. Options: recycling (needs sorting and downgrades quality), incineration (energy recovery, but toxic gases from PVC), and developing biodegradable or photodegradable polymers. None is a complete answer.

The bit that catches people outCondensation polymers can be hydrolysed back to their monomers because the ester and amide links are attackable; addition polymers cannot, because their backbone is unreactive C–C. That single structural fact is why plastic waste behaves so differently from natural polymers.

The grown-up words

What it means	What it is called	Note
Polymer formed with nothing eliminated	addition polymer	From alkene monomers
Polymer formed losing a small molecule	condensation polymer	Water or HCl
Smallest repeating section of a polymer	repeat unit	Bonds through the brackets
Polymer joined by ester links	polyester	From a diol and a diacid
Polymer joined by amide links	polyamide	Nylon, Kevlar, proteins
Additive pushing chains apart	plasticiser	Increases flexibility

Check you have got it

Why can nylon be hydrolysed but poly(ethene) cannot?

Nylon's chain contains amide links which can be attacked by water under acidic or alkaline conditions. Poly(ethene) has only unreactive C–C and C–H bonds in its backbone.

Give one advantage and one problem with incinerating waste plastics.

Advantage: the energy released can generate electricity, and the waste volume is greatly reduced. Problem: burning PVC releases HCl and other toxic gases, requiring expensive scrubbing.

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Topic 21

Organic synthesis (AS)

Synthesis is planning a route: getting from a starting material to a target in as few reliable steps as possible.

Picture itPlanning a journey with a bus map. You cannot fly directly, so you look for connections. Every reaction you have learned is one bus route between two functional groups, and synthesis is finding the shortest sequence of routes that connects where you are to where you need to be.

Think in functional group conversions

Alkane → halogenoalkane (UV, X₂). Halogenoalkane → alcohol (aqueous NaOH), → nitrile (ethanolic KCN), → amine (ethanolic NH₃), → alkene (hot ethanolic NaOH). Alcohol → aldehyde/ketone/acid (oxidation), → alkene (dehydration). Alkene → alcohol, halogenoalkane, dihalide.

Changing the carbon chain length

Only a few reactions add a carbon: halogenoalkane + CN⁻, and carbonyl + HCN. Both give a nitrile that can be hydrolysed to an acid or reduced to an amine. If the target has more carbons than the start, one of these is almost certainly in the route.

Working backwards

Start from the target and ask what could make it in one step; then ask the same of that intermediate. This retrosynthesis converges on the starting material far faster than guessing forwards.

Reagents and conditions carry the marks

Naming the reagent alone is rarely enough — aqueous versus ethanolic, reflux versus distil, hot versus cold, catalyst, temperature. This is where most synthesis marks are actually won and lost.

Practical realities

Reflux for heating without loss, distillation for separating by boiling point, separating funnel for immiscible layers, recrystallisation for purification, and melting point as a purity check — sharp and matching for a pure sample.

The bit that catches people outThe shortest route on paper is not always the best. A step with a poor yield, a mixture of products, or a hazardous reagent may make a longer route preferable — and saying why you chose your route is often worth as much as the route itself.

The grown-up words

What it means	What it is called	Note
Planning backwards from the target	retrosynthesis	Converges quickly
Heating without losing volatile material	reflux	Condenser vertical
Separating liquids by boiling point	distillation	Collect at a set temperature
Purifying a solid by dissolving and cooling	recrystallisation	Impurities stay in solution
Sharp value confirming purity	melting point	Broad and low means impure
Reaction adding a carbon via cyanide	nitrile formation	Then hydrolyse or reduce

Check you have got it

How would you convert bromoethane into propanoic acid?

Reflux with ethanolic KCN to give propanenitrile (adding a carbon), then hydrolyse with dilute acid under reflux to give propanoic acid.

Why might a two-step route be preferred over a one-step route with the same product?

The one-step route may have a poor yield, produce a mixture requiring difficult separation, or use a hazardous reagent — so the longer route can be cheaper, safer or purer overall.

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Topic 22

Analytical techniques (AS)

Deducing an unknown structure from spectra is detective work, and each technique answers a different question.

Picture itThree witnesses to the same event. Mass spectrometry tells you how heavy the molecule is and how it breaks. Infrared tells you which functional groups are present. NMR tells you how the hydrogens are arranged. No one witness gives the whole story; together they usually do.

Mass spectrometry

The molecular ion peak at the highest m/e gives the relative molecular mass. Fragmentation patterns give structural clues — a loss of 15 suggests CH₃, 29 suggests CHO or C₂H₅, 17 suggests OH. The M+1 peak comes from carbon-13 and its size indicates the number of carbons.

Isotope patterns are a giveaway

Chlorine gives M and M+2 in roughly 3:1 because of ³⁵Cl and ³⁷Cl. Bromine gives M and M+2 in roughly 1:1. Spotting that pattern identifies the halogen immediately.

Infrared spectroscopy

Bonds absorb at characteristic wavenumbers: O–H in alcohols is broad around 3200–3600 cm⁻¹, O–H in acids is very broad around 2500–3300, C=O is sharp and strong around 1700, N–H around 3300–3500. The fingerprint region below 1500 cm⁻¹ identifies a specific compound by comparison.

Reading an IR spectrum

Look only for the key peaks and say what each shows — a broad peak at 3000 and a sharp one at 1700 together mean a carboxylic acid. Trying to assign every peak wastes time and earns nothing.

Carbon-13 NMR

The number of peaks equals the number of chemically different carbon environments. Chemical shifts, read from the data sheet, indicate the environment. Symmetry reduces the peak count — which is often the decisive clue between two candidate structures.

The bit that catches people outIn infrared, a broad O–H peak around 3300 cm⁻¹ is an alcohol, but a very broad one stretching from about 2500 to 3300 and overlapping the C–H peaks is a carboxylic acid. The difference is in the width, and it is the difference between two answers.

The grown-up words

What it means	What it is called	Note
Peak at the highest m/e giving Mr	molecular ion peak	The intact molecule
Breaking of the molecular ion	fragmentation	Gives structural clues
Peak from carbon-13 next to the molecular ion	M+1 peak	Indicates carbon count
M and M+2 in 3:1 ratio	chlorine isotope pattern	Bromine gives 1:1
Region below 1500 cm-1 unique to a compound	fingerprint region	Matched against a database
Number of chemically distinct carbons	carbon environments	Equals the peak count

Check you have got it

A spectrum shows a strong sharp peak at 1715 cm⁻¹ and no broad peak above 3000. What does this suggest?

A C=O group with no O–H — so an aldehyde, ketone or ester, but not a carboxylic acid or alcohol.

Propan-2-ol gives two peaks in its carbon-13 NMR spectrum. Why not three?

The two CH₃ groups are in identical chemical environments because of the molecule's symmetry, so they produce a single peak.

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Topic 23

Chemical energetics (A2)

A reaction happens if it lowers the total disorder-adjusted energy — and that is a different question from whether it happens quickly.

Picture itA ball at the top of a hill will roll down: the process is feasible. Whether it actually moves depends on whether someone gives it a nudge over the small ridge in front of it. Thermodynamics tells you about the hill; kinetics tells you about the ridge. Both matter, and they are separate questions.

Entropy is a measure of dispersal

Entropy (S) increases with disorder and with the number of ways particles and energy can be arranged. Gases have far higher entropy than liquids, which have more than solids. More moles of gas on the product side means a positive ΔS^⦵.

Calculating entropy change

ΔS^⦵ = ΣS^⦵(products) − ΣS^⦵(reactants), in J K⁻¹ mol⁻¹. Note the units: entropy is in joules, enthalpy in kilojoules, so one of them must be converted before combining them.

Gibbs free energy decides feasibility

ΔG^⦵ = ΔH^⦵ − TΔS^⦵. A reaction is feasible when ΔG is negative. Because of the T term, a reaction with positive ΔH and positive ΔS becomes feasible above a certain temperature — set ΔG = 0 and T = ΔH ÷ ΔS gives that temperature.

The four sign combinations

ΔH negative, ΔS positive: always feasible. ΔH positive, ΔS negative: never. ΔH negative, ΔS negative: feasible at low T. ΔH positive, ΔS positive: feasible at high T. Being able to reason through all four is a standard question.

Lattice energy and Born–Haber revisited

Lattice energy becomes more exothermic with higher ionic charge and smaller ionic radius. Comparing the experimental value from a Born–Haber cycle with the theoretical value assuming a purely ionic model reveals covalent character — a bigger discrepancy means more covalent character, as predicted by Fajans' rules.

The bit that catches people outFeasible is not the same as fast. The conversion of diamond to graphite has a negative ΔG at room temperature and is thermodynamically feasible — the activation energy is so enormous that it never happens. Never answer a ΔG question with a rate argument.

The grown-up words

What it means	What it is called	Note
Measure of dispersal of energy and matter	entropy	Units J per K per mol
dH minus T dS	Gibbs free energy change	Negative means feasible
Reaction that can occur thermodynamically	feasible reaction	Says nothing about rate
Temperature at which dG becomes zero	minimum feasible temperature	T = dH / dS
Energy released forming a lattice from gaseous ions	lattice energy	More exothermic if small and highly charged
Difference between experimental and theoretical lattice energy	covalent character	Predicted by Fajans' rules

Check you have got it

A reaction has ΔH = +120 kJ mol⁻¹ and ΔS = +200 J K⁻¹ mol⁻¹. Above what temperature is it feasible?

T = ΔH ÷ ΔS = 120 000 ÷ 200 = 600 K. Note the conversion of kJ to J.

Why is a reaction with negative ΔH and negative ΔS feasible only at low temperatures?

ΔG = ΔH − TΔS. With ΔS negative, −TΔS is positive and grows with temperature, so at high T it outweighs the negative ΔH and ΔG becomes positive.

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Topic 24

Electrochemistry (A2)

Standard electrode potentials rank every half reaction on one scale, so you can predict any redox reaction before performing it.

Picture itA league table. Once every half cell has a number measured against the same reference, you can look up any two and know which will oxidise which — without ever mixing them. That is what the standard hydrogen electrode gives you.

The standard hydrogen electrode

Defined as exactly 0.00 V under standard conditions: 298 K, 100 kPa H₂, 1.00 mol dm⁻³ H⁺, platinum electrode. Every other potential is measured relative to it, which is what makes the scale meaningful.

Reading E^⦵ values

A more positive E^⦵ means a greater tendency to be reduced — a better oxidising agent. E^⦵_cell = E^⦵(more positive) − E^⦵(less positive), and a positive cell potential means the reaction is feasible. The species with the more positive value is reduced; the other is oxidised.

Cell diagrams and conventions

Written with the more negative electrode on the left: Zn | Zn²⁺ ‖ Cu²⁺ | Cu. A single line is a phase boundary, a double line is the salt bridge. Getting the convention right is worth marks in its own right.

The Nernst equation

E = E^⦵ + (RT/zF) ln([oxidised]/[reduced]). In practice: increasing the concentration of the oxidised species makes E more positive. This explains why real cell potentials drift from the tabulated values as a cell discharges.

Applications

Rechargeable cells, hydrogen–oxygen fuel cells (high efficiency, water as the only product, but hydrogen storage and production remain difficult), electrolysis quantities via Q = It and F, and predicting whether a proposed redox reaction will actually work.

The bit that catches people outA positive E^⦵_cell means feasible, not fast — the same distinction as ΔG. It also assumes standard conditions; changing concentrations shifts the value, and a reaction that is barely feasible on paper may not proceed in practice.

The grown-up words

What it means	What it is called	Note
Reference electrode defined as 0.00 V	standard hydrogen electrode	298 K, 100 kPa, 1 mol dm-3
Tendency of a half cell to be reduced	standard electrode potential	More positive = better oxidant
Difference between two electrode potentials	cell potential	Positive means feasible
Connection allowing ion flow between half cells	salt bridge	Double line in a cell diagram
Equation relating E to concentration	Nernst equation	Explains non-standard values
Cell converting fuel directly to electricity	fuel cell	Water is the only product

Check you have got it

Given E⦵(Cu²⁺/Cu) = +0.34 V and E⦵(Zn²⁺/Zn) = −0.76 V, what is E⦵cell and which is oxidised?

E⦵cell = 0.34 − (−0.76) = +1.10 V. Zinc has the more negative value so it is oxidised; copper(II) is reduced.

Why does the voltage of a cell fall as it is used?

The concentration of the oxidised species falls and the reduced species rises, which by the Nernst equation reduces the electrode potential and hence the cell potential.

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Topic 25

Equilibria (A2)

Acid-base chemistry becomes quantitative here — you can calculate pH at any point, and design a solution that resists change.

Picture itA shock absorber for pH. A buffer contains both a weak acid and its conjugate base in reasonable quantity, so added acid is mopped up by the base and added alkali by the acid. Blood is buffered at pH 7.35–7.45, and a shift of a few tenths is a medical emergency — which is the measure of how well it works.

K_a, pK_a and pH calculations

For a weak acid, K_a = [H⁺]²/[HA] approximately, so [H⁺] = √(K_a × [HA]). pK_a = −log K_a, and a smaller pK_a means a stronger acid. For a strong acid, [H⁺] simply equals the concentration.

Buffers and the Henderson–Hasselbalch relationship

pH = pK_a + log([salt]/[acid]). A buffer works best when [salt] = [acid], giving pH = pK_a. It resists change because both components are present in large amounts relative to what is added.

Titration curves have four regions to read

The initial pH, the buffer region where pH changes slowly, the steep vertical section at the equivalence point, and the plateau beyond. Strong–strong gives equivalence at pH 7; weak acid–strong base above 7; strong acid–weak base below 7.

Choosing an indicator

The indicator's pK_in must lie within the vertical section of the curve. Methyl orange (3.1–4.4) suits strong acid–weak base; phenolphthalein (8.3–10.0) suits weak acid–strong base. A weak acid–weak base titration has no vertical section, so no indicator works.

Solubility product

K_sp = [A⁺]^a[B⁻]^b for a sparingly soluble salt. A precipitate forms when the ionic product exceeds K_sp. The common ion effect reduces solubility: adding a shared ion pushes the equilibrium back towards the solid.

The bit that catches people outA buffer does not hold pH constant — it changes the pH far less than the same addition would to unbuffered water. Saying a buffer 'keeps the pH the same' is not quite right, and the mark scheme wants 'resists change'.

The grown-up words

What it means	What it is called	Note
Acid dissociation constant	Ka	Larger means stronger acid
Negative log of Ka	pKa	Smaller means stronger acid
Solution resisting pH change	buffer	Weak acid plus its conjugate base
Point where moles of acid and base are equal	equivalence point	Not always pH 7
Product of ion concentrations at saturation	solubility product	Ksp
Reduced solubility from a shared ion	common ion effect	Shifts equilibrium to the solid

Check you have got it

Calculate the pH of 0.10 mol dm⁻³ ethanoic acid, Ka = 1.8 × 10⁻⁵.

[H⁺] = √(1.8 × 10⁻⁵ × 0.10) = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³. pH = −log(1.34 × 10⁻³) = 2.87.

Why is phenolphthalein unsuitable for a strong acid–weak base titration?

Its colour change (pH 8.3–10.0) lies outside the vertical section of that curve, which occurs below pH 7, so the colour would change well before the equivalence point.

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Topic 26

Reaction kinetics (A2)

Rate equations are found by experiment, and once you have one it tells you something about the mechanism that no amount of theory could.

Picture itYou cannot see a mechanism. But if doubling the concentration of one reactant doubles the rate and doubling another does nothing, you have learned that the second one is not involved in the slowest step. Rate equations are an indirect window onto events too fast to observe.

Rate equations come from data, never from the equation

rate = k[A]^m[B]ⁿ. The orders m and n must be determined experimentally — they bear no necessary relation to the stoichiometric coefficients. The overall order is m + n, and the units of k depend on the overall order.

Determining order

Initial rates: change one concentration at a time and see how the rate responds. Graphs: a zero-order plot of concentration against time is a straight line; first-order has a constant half-life; second-order gives a straight line when 1/[A] is plotted against time.

Rate-determining step

The slowest step governs the overall rate, and only species involved in it (or before it) appear in the rate equation. A proposed mechanism must be consistent with the experimental rate equation — that consistency is the evidence for it.

SN1 and SN2 as the worked example

SN1 is first order overall, because the rate-determining step is the halogenoalkane forming a carbocation alone. SN2 is second order, because both the halogenoalkane and the nucleophile are in the single step. The kinetics distinguish the two mechanisms.

The Arrhenius equation

k = Ae^−Ea/RT, or ln k = ln A − E_a/RT. Plotting ln k against 1/T gives a straight line of gradient −E_a/R, which is how activation energies are measured. Also: homogeneous catalysis works through an intermediate; heterogeneous catalysis through adsorption onto a surface.

The bit that catches people outYou cannot read the order off the balanced equation. A reaction with 2A in the equation may be first order in A, or zero order. The orders come from experiment, and writing them from the stoichiometry is a standard and heavily penalised mistake.

The grown-up words

What it means	What it is called	Note
Equation relating rate to concentrations	rate equation	Found experimentally
Power to which a concentration is raised	order of reaction	Not from the stoichiometry
Slowest step governing the overall rate	rate-determining step	Species in it appear in the rate equation
Time for concentration to halve	half-life	Constant for first order
Constant in the rate equation	rate constant	Units depend on overall order
ln k against 1/T gives -Ea/R	Arrhenius equation	How Ea is measured

Check you have got it

Doubling [A] quadruples the rate; doubling [B] leaves it unchanged. Write the rate equation.

rate = k[A]². The reaction is second order in A and zero order in B, so B does not appear.

A reaction shows a constant half-life. What order is it, and why does that follow?

First order. Only for first order is the half-life independent of the starting concentration, because the rate is directly proportional to the concentration remaining.

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Topic 27

Group 2 (A2)

The group 2 trends you learned at AS have quantitative explanations at A2, and they come down to charge density.

Picture itA small, heavily charged ion is like a strong hand gripping a balloon — it distorts what it holds. A big, weakly charged ion barely dents it. Everything in this topic — thermal stability, solubility, hydration — follows from how hard the cation pulls on its neighbours.

Charge density explains the decomposition trend

Going down group 2, cations get larger, so charge density falls. A small, highly charged cation like Mg²⁺ polarises the carbonate ion strongly, weakening a C–O bond and making decomposition easier. Ba²⁺ polarises much less, so BaCO₃ decomposes at a much higher temperature.

Solubility trends have a thermodynamic explanation

Dissolving involves breaking the lattice (endothermic, governed by lattice energy) and hydrating the ions (exothermic, governed by hydration enthalpy). Both fall down the group but at different rates — which is why hydroxide solubility rises while sulfate solubility falls.

Hydration enthalpy

The energy released when one mole of gaseous ions is hydrated. It becomes less exothermic as ions get larger and as charge falls — the same charge density argument, applied to the ion's interaction with water rather than with the lattice.

Enthalpy of solution

ΔH_sol = −lattice energy + Σ hydration enthalpies. Whether a salt dissolves depends on the balance, and because the two terms are large and of opposite sign, small changes in either can flip the result.

Applying the reasoning

Questions in this topic rarely ask you to recall a fact — they ask you to explain a trend using ionic radius, charge and their effect on polarising power or hydration. The structure of the answer is almost always the same, and it is worth rehearsing.

The bit that catches people outBoth lattice energy and hydration enthalpy become less exothermic down a group. Solubility trends depend on which falls faster, which is why hydroxides and sulfates trend in opposite directions. Quoting only one of the two terms cannot explain either trend.

The grown-up words

What it means	What it is called	Note
Charge divided by ionic radius	charge density	High for small, highly charged ions
Ability to distort a neighbouring ion	polarising power	Greatest for small cations
Energy released hydrating one mole of gaseous ions	hydration enthalpy	Less exothermic for larger ions
Energy released forming a lattice from gaseous ions	lattice energy	Less exothermic down a group
Sum of lattice breaking and hydration	enthalpy of solution	Small difference of large numbers
Breaking down on heating	thermal decomposition	Harder down group 2

Check you have got it

Explain, using charge density, why MgCO₃ decomposes at a lower temperature than CaCO₃.

Mg²⁺ is smaller than Ca²⁺, so it has a higher charge density and polarises the carbonate ion more strongly. This weakens a C–O bond, so less energy is needed to break the carbonate into the oxide and CO₂.

Why can solubility not be predicted from lattice energy alone?

Dissolving requires breaking the lattice (endothermic) and hydrating the ions (exothermic). Solubility depends on the balance between the two, so hydration enthalpy must be considered as well.

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Topic 28

Transition elements

Partly filled d orbitals give the transition metals four properties nothing else in the periodic table has.

Picture itA set of orbitals almost exactly equal in energy — until a ligand approaches and splits them by a gap that happens to match visible light. That single accident of energy scale is why transition metal compounds are coloured, and why so much of biology and industry depends on them.

The defining property

A transition element forms at least one stable ion with a partly filled d subshell. Scandium and zinc are excluded — Sc³⁺ has an empty d subshell and Zn²⁺ a full one, so neither shows the characteristic properties.

Variable oxidation states

4s and 3d electrons are similar in energy, so several can be removed with comparable ease. Manganese ranges from +2 to +7. This is what makes transition metals such versatile catalysts — they can gain and lose electrons reversibly.

Complex ions and ligands

A ligand is a species with a lone pair that forms a dative bond to the central ion. Monodentate (H₂O, NH₃, Cl⁻), bidentate (ethanedioate, en) and multidentate (EDTA). Six-coordinate complexes are octahedral, four-coordinate are tetrahedral or square planar.

Colour and d-orbital splitting

Ligands split the five d orbitals into two energy levels. An electron absorbs a photon of exactly the splitting energy and jumps up; the colour you see is the complement of the light absorbed. Changing the ligand changes the splitting, and so changes the colour.

Ligand exchange and stability constants

Ligands can be replaced, sometimes with a colour and shape change — [Cu(H₂O)₆]²⁺ blue to [CuCl₄]²⁻ yellow. The stability constant K_stab measures how favourable the exchange is; a larger value means a more stable complex, which is why EDTA displaces almost everything.

The bit that catches people outA compound's colour is the light it does not absorb. Copper(II) sulfate solution looks blue because it absorbs red-orange light; the blue is what passes through. Saying the complex 'absorbs blue light' reverses the whole explanation.

The grown-up words

What it means	What it is called	Note
Element forming an ion with a partly full d subshell	transition element	Excludes Sc and Zn
Lone pair donor bonded to a central metal ion	ligand	Forms a dative bond
Number of dative bonds to the central ion	coordination number	Usually 6 or 4
Ligand forming two bonds to the metal	bidentate ligand	e.g. ethanedioate
Difference in energy between split d orbitals	d-orbital splitting	Matches visible light
Replacement of one ligand by another	ligand exchange	Often changes colour
Equilibrium constant for complex formation	stability constant	Larger = more stable

Check you have got it

Why is zinc not classed as a transition element?

Its only common ion, Zn²⁺, has a completely full 3d subshell, so it does not show variable oxidation states, coloured compounds or the catalytic behaviour that define the transition elements.

Explain why [Cu(H₂O)₆]²⁺ is pale blue but [Cu(NH₃)₄(H₂O)₂]²⁺ is deep blue.

Ammonia is a stronger-field ligand than water, so it splits the d orbitals by a larger energy gap. A different wavelength is absorbed, so a different colour is transmitted.

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Topic 29

An introduction to A Level organic chemistry

The A2 organic course adds mechanisms that must be drawn precisely and stereochemistry that must be described exactly.

Picture itAt AS you learned which reactions happen. At A2 you have to show the electrons moving, and say what the three-dimensional result is. The chemistry is the same; the level of proof demanded is higher.

Curly arrow conventions, strictly

A double-headed arrow shows a pair of electrons moving; a single-headed 'fishhook' shows one electron in radical mechanisms. Arrows begin at a bond or a lone pair — never at an atom, never at a positive charge — and end where the pair goes.

The mechanism types you must know

Free radical substitution, electrophilic addition, electrophilic substitution (arenes), nucleophilic substitution (SN1 and SN2), nucleophilic addition (carbonyls), nucleophilic addition–elimination (acyl chlorides), elimination, and condensation.

Stereochemistry in mechanisms

SN2 proceeds with inversion of configuration, because the nucleophile attacks from the opposite side to the leaving group. SN1 goes through a planar carbocation, so attack is equally likely from either face and a racemic mixture results — a difference you can detect optically.

Optical activity

Enantiomers rotate plane-polarised light equally in opposite directions, measured with a polarimeter. A racemic mixture is optically inactive because the rotations cancel. Enantiomers are otherwise identical in every physical property, which is what makes them hard to separate.

Why stereochemistry matters practically

Biological receptors are themselves chiral, so two enantiomers of a drug can have entirely different effects. Producing a single enantiomer rather than a racemic mixture is a major concern in pharmaceutical synthesis.

The bit that catches people outSN1 gives a racemic mixture and SN2 gives inversion. That difference is not a detail — it is experimental evidence distinguishing the two mechanisms, alongside the kinetics, and questions often ask you to use it as such.

The grown-up words

What it means	What it is called	Note
Double-headed arrow showing a pair moving	curly arrow	Starts at a bond or lone pair
Single-headed arrow for one electron	fishhook arrow	Radical mechanisms
Nucleophile attacks opposite the leaving group	inversion of configuration	Characteristic of SN2
Equal mixture of two enantiomers	racemic mixture	Optically inactive
Non-superimposable mirror image forms	enantiomers	Rotate light oppositely
Instrument measuring optical rotation	polarimeter	Detects chirality

Check you have got it

Why does an SN1 reaction on a chiral halogenoalkane give a racemic product?

The carbocation intermediate is planar, so the nucleophile is equally likely to attack from either face, giving equal amounts of both enantiomers.

Why must an arrow never start at a positive charge?

Curly arrows show the movement of electrons, and a positive charge indicates a shortage of electrons — there is nothing there to move. The arrow must start at the bond or lone pair supplying the electrons.

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Topic 30

Hydrocarbons: arenes

Benzene's delocalised ring makes it far more stable than its structure suggests, and that stability dictates how it reacts.

Picture itBenzene should behave like three alkenes joined in a ring. It does not. It refuses to add, and instead substitutes — because addition would destroy a delocalised system worth about 150 kJ mol⁻¹ of extra stability. Everything in this topic is benzene protecting that ring.

Evidence for delocalisation

All six C–C bonds are the same length, between a single and a double bond. The enthalpy of hydrogenation is about 150 kJ mol⁻¹ less exothermic than three times that of cyclohexene. Benzene resists addition reactions that alkenes undergo readily.

Electrophilic substitution

The ring's π system attracts an electrophile, forming an unstable intermediate, which then loses H⁺ to restore the delocalised ring. Restoring aromaticity is the driving force, and it is why substitution wins over addition.

The four reactions

Nitration: concentrated HNO₃ and H₂SO₄ at about 55 °C, generating NO₂⁺. Halogenation: Cl₂ or Br₂ with AlCl₃ or FeBr₃ as a halogen carrier. Friedel–Crafts alkylation: RCl with AlCl₃. Friedel–Crafts acylation: RCOCl with AlCl₃, giving a ketone.

Side chain versus ring

With UV light and no halogen carrier, a methyl side chain undergoes free radical substitution. With a halogen carrier and no UV, the ring is substituted. Same reagents, opposite site of attack — decided entirely by conditions.

Phenol is not an alcohol

The OH's lone pair delocalises into the ring, making phenol acidic enough to react with NaOH (but not with carbonates), and activating the ring so that bromination happens with bromine water alone, giving 2,4,6-tribromophenol as a white precipitate.

The bit that catches people outPhenol is more acidic than an alcohol and more reactive to substitution than benzene, and both facts come from the same cause: the OH lone pair delocalises into the ring. One structural feature, two consequences that seem unrelated until you see the link.

The grown-up words

What it means	What it is called	Note
Ring of delocalised pi electrons	aromatic system	Extra stability
Extra stability from delocalisation	delocalisation energy	About 150 kJ per mole
Electrophile replaces a hydrogen on the ring	electrophilic substitution	Restores aromaticity
Catalyst generating the electrophile from a halogen	halogen carrier	AlCl3 or FeBr3
Attaching an alkyl or acyl group to a ring	Friedel-Crafts reaction	Needs AlCl3
OH group attached directly to a benzene ring	phenol	Weakly acidic

Check you have got it

Why does benzene undergo substitution rather than addition?

Addition would break the delocalised π system and lose about 150 kJ mol⁻¹ of stabilisation. Substitution restores the aromatic ring, so it is energetically far more favourable.

Why does phenol react with bromine water without a catalyst, while benzene does not?

The oxygen lone pair is delocalised into the ring, increasing its electron density and making it far more attractive to electrophiles, so no halogen carrier is needed.

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Topic 31

Halogen compounds (A2)

At A2 the halogenoalkane reactions you know are backed by kinetic evidence and stereochemical proof.

Picture itYou already know that halogenoalkanes react with nucleophiles. The A2 question is how — and the answer is settled not by argument but by two independent lines of evidence: what the rate equation says, and what happens to the optical activity.

SN1 and SN2 compared properly

SN2: one step, rate = k[RX][Nu], inversion of configuration, favoured by primary halogenoalkanes where the carbon is accessible. SN1: two steps via a carbocation, rate = k[RX] only, racemisation, favoured by tertiary where the carbocation is stabilised.

Why substitution pattern decides the mechanism

Tertiary carbocations are stabilised by three electron-releasing alkyl groups, so SN1 is accessible. Primary carbons have little steric hindrance, so backside attack is easy and SN2 dominates. Secondary halogenoalkanes sit between the two and can show both.

Elimination versus substitution revisited

Strong base, high temperature and ethanolic conditions favour elimination; a good nucleophile in aqueous conditions favours substitution. Bulky bases also push towards elimination because attack at the carbon is hindered.

Aryl halides are unreactive

In chlorobenzene the halogen's lone pair overlaps with the ring's π system, giving partial double bond character and shortening the C–Cl bond. It is far harder to break, so aryl halides resist the nucleophilic substitution that halogenoalkanes undergo easily.

Uses and consequences

Halogenoalkanes remain important as synthetic intermediates and solvents. The CFC story is the standard case study: a compound chosen precisely for its inertness turned out to be inert enough to reach the stratosphere intact.

The bit that catches people outChlorobenzene and chloroethane look similar and behave completely differently. The lone pair delocalisation into the ring is the whole reason — and it is a good example of why you cannot reason about reactivity from the formula alone.

The grown-up words

What it means	What it is called	Note
One-step substitution, second order overall	SN2	Inversion, primary favoured
Two-step substitution via a carbocation	SN1	Racemisation, tertiary favoured
Bulk of groups blocking attack	steric hindrance	Pushes towards SN1 or elimination
Halogen attached directly to a benzene ring	aryl halide	Very unreactive
Partial double bond from lone pair overlap	delocalisation	Strengthens the C-X bond
Loss of optical activity in a product	racemisation	Evidence for SN1

Check you have got it

Predict the mechanism for 2-bromo-2-methylpropane with aqueous NaOH, with a reason.

SN1. It is tertiary, so the carbocation formed is stabilised by three electron-releasing alkyl groups, and backside attack is sterically blocked.

Why is chlorobenzene not hydrolysed by aqueous NaOH under normal conditions?

The chlorine's lone pair overlaps with the ring's delocalised system, giving the C–Cl bond partial double bond character. It is shorter and stronger, so it does not break readily.

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Topic 32

Hydroxy compounds (A2)

Phenols look like alcohols and are not, and the difference in one bond explains every difference in behaviour.

Picture itEthanol and phenol both end in OH. Ethanol is neutral and unreactive to bromine water. Phenol is acidic and reacts with bromine water instantly, giving a white precipitate. Attaching that OH to a benzene ring changes it completely, because the ring pulls the lone pair in.

Why phenol is acidic

The oxygen lone pair delocalises into the ring, which weakens the O–H bond and stabilises the phenoxide ion formed on dissociation. Phenol therefore reacts with NaOH to give sodium phenoxide, but is too weak to react with sodium carbonate — which is exactly how you distinguish it from a carboxylic acid.

Why the ring is activated

The same delocalisation increases electron density in the ring, so phenol undergoes electrophilic substitution far more readily than benzene. Bromine water alone gives 2,4,6-tribromophenol as a white precipitate — no halogen carrier is needed and the reaction is immediate.

Reactions of phenol

With NaOH → sodium phenoxide and water. With bromine water → 2,4,6-tribromophenol. With acyl chlorides → esters (phenol is too weak a nucleophile for direct esterification with a carboxylic acid). With neutral FeCl₃ → a violet colour, a classic test.

Comparing alcohols, phenols and carboxylic acids

Reaction with sodium: all three. Reaction with NaOH: phenol and acid, not alcohol. Reaction with Na₂CO₃: acid only. That three-test sequence unambiguously distinguishes the three classes and is a standard practical question.

Uses

Phenols are used in antiseptics and disinfectants, in resins such as Bakelite, and as intermediates in making polymers, dyes and pharmaceuticals — most of which exploit the activated ring.

The bit that catches people outPhenol reacts with NaOH but not with Na₂CO₃; a carboxylic acid reacts with both. That single difference is the definitive test between them, and it follows directly from phenol being the weaker acid.

The grown-up words

What it means	What it is called	Note
OH attached directly to a benzene ring	phenol	Weakly acidic
Ion formed when phenol loses H+	phenoxide ion	Stabilised by delocalisation
Ring made more reactive to electrophiles	activated ring	By the OH lone pair
White precipitate with bromine water	2,4,6-tribromophenol	Test for phenol
Violet colour with neutral iron(III) chloride	phenol test	Classic identification
Acid reacting with carbonates	carboxylic acid	Phenol does not

Check you have got it

Explain why phenol is a stronger acid than ethanol.

The negative charge on the phenoxide ion is delocalised into the benzene ring, stabilising it. The ethoxide ion has no such stabilisation, so ethanol dissociates far less.

How can you distinguish phenol from ethanoic acid in the laboratory?

Add sodium carbonate solution. Ethanoic acid effervesces, releasing CO₂; phenol does not react because it is too weak an acid.

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Topic 33

Carboxylic acids and derivatives (A2)

The derivatives of carboxylic acids differ enormously in reactivity, and the order has one clear explanation.

Picture itA ladder of reactivity. Acyl chlorides react violently with water; esters need heating with a catalyst; amides need prolonged reflux with concentrated acid. The order runs acyl chloride > ester > amide, and it tracks how good the leaving group is.

Relative reactivity and why

The better the leaving group, the more reactive the derivative. Chloride is an excellent leaving group, so acyl chlorides are extremely reactive. Alkoxide is worse, so esters are moderate. Amide ion is very poor, so amides are the least reactive and need forcing conditions.

Nucleophilic addition–elimination

The nucleophile adds to the δ+ carbonyl carbon to give a tetrahedral intermediate, which then eliminates the leaving group and reforms the C=O. This one mechanism covers the reactions of acyl chlorides with water, alcohols, ammonia and amines.

Acidity and substituent effects

Electron-withdrawing groups near the COOH stabilise the carboxylate ion and increase acidity — so chloroethanoic acid is a stronger acid than ethanoic acid, and trichloroethanoic acid stronger still. Electron-releasing alkyl groups have the opposite effect.

Hydrolysis, acid and alkaline

Acid hydrolysis of an ester is reversible, giving the carboxylic acid and the alcohol. Alkaline hydrolysis is irreversible, giving the carboxylate salt, because that ion does not react back with the alcohol. Amide hydrolysis gives the acid plus ammonium salt (acid conditions) or the carboxylate plus ammonia (alkaline).

Applications

Esters as flavourings, solvents and plasticisers; polyesters as fibres; acyl chlorides as reactive intermediates in synthesis; amides as the link in nylon and in proteins. Aspirin is made by acylating salicylic acid with ethanoic anhydride.

The bit that catches people outAmides are neutral, not basic, even though they contain nitrogen. The lone pair is delocalised into the adjacent carbonyl group and is not available to accept a proton — which is why an amide behaves nothing like an amine.

The grown-up words

What it means	What it is called	Note
Highly reactive RCOCl compound	acyl chloride	Chloride is a good leaving group
Add then eliminate at a carbonyl	nucleophilic addition-elimination	Via a tetrahedral intermediate
Species departing with the bonding pair	leaving group	Better group means faster reaction
Group pulling electron density away	electron-withdrawing group	Increases acid strength
Irreversible ester hydrolysis with alkali	saponification	Gives the carboxylate salt
Neutral nitrogen compound with a C=O	amide	Lone pair delocalised

Check you have got it

Why is chloroethanoic acid a stronger acid than ethanoic acid?

The electronegative chlorine withdraws electron density, spreading and stabilising the negative charge on the carboxylate ion, so dissociation is more favourable.

Why do acyl chlorides react with water far more readily than esters?

Chloride is a much better leaving group than alkoxide, so the elimination step of the addition–elimination mechanism is far easier.

Edvia Free Resources · Chemistry 9701 · Topic 33 — free to copy and share

Topic 34

Nitrogen compounds (A2)

Amines, amides and amino acids all contain nitrogen and behave in three different ways, depending on what the lone pair is doing.

Picture itThe nitrogen lone pair is the whole story. Free and pushed forward by alkyl groups, it is a good base — an amine. Delocalised into a carbonyl, it is unavailable — an amide, neutral. Delocalised into a benzene ring, it is partly withdrawn — phenylamine, a weak base.

Base strength compared

Aliphatic amines > ammonia > aromatic amines > amides. Alkyl groups release electrons and make the lone pair more available; a benzene ring or a carbonyl delocalises it away. Explaining the order using lone pair availability is the standard question.

Preparing amines

Halogenoalkane with excess ethanolic ammonia (mixture of products). Nitrile reduction with LiAlH₄ or H₂/Ni (clean, extends the chain). Nitrobenzene reduction with Sn and concentrated HCl, then NaOH, gives phenylamine.

Diazotisation and azo dyes

Phenylamine with nitrous acid (from NaNO₂ and HCl) below 10 °C gives a diazonium salt, which couples with phenol or another aromatic amine to give a highly coloured azo compound. The extended delocalised system across the N=N link is what produces the colour.

Amino acids and proteins

Amino acids exist as zwitterions; the isoelectric point is where there is no net charge. Peptide bonds form by condensation. Electrophoresis separates amino acids by charge at a chosen pH — those below their isoelectric point are positive and move to the cathode.

Protein structure

Primary, secondary (α-helix and β-pleated sheet, held by hydrogen bonds), tertiary (hydrogen bonds, ionic bonds, disulfide bridges, hydrophobic interactions) and quaternary. Hydrolysis with 6 mol dm⁻³ HCl breaks a protein into its amino acids for analysis.

The bit that catches people outDiazotisation must be done below 10 °C. Above that, the diazonium salt decomposes and the coupling fails — a rare case where a specific temperature is itself the examinable point.

The grown-up words

What it means	What it is called	Note
Compound with an available nitrogen lone pair	amine	Acts as a base
Neutral nitrogen compound next to a C=O	amide	Lone pair delocalised
Salt made from an aromatic amine below 10 C	diazonium salt	Decomposes if warmed
Coloured compound with an N=N link	azo compound	Used as dyes
Separating molecules by charge in a field	electrophoresis	Depends on the pH used
Molecule with both positive and negative charge	zwitterion	Amino acids in solution

Check you have got it

Why is ethylamine a stronger base than phenylamine?

In ethylamine the ethyl group releases electron density towards nitrogen, making the lone pair more available to accept a proton. In phenylamine the lone pair is delocalised into the ring and is much less available.

At a pH below its isoelectric point, which electrode does an amino acid move towards?

The cathode (negative electrode), because at low pH the amino group is protonated and the molecule carries a net positive charge.

Edvia Free Resources · Chemistry 9701 · Topic 34 — free to copy and share

Topic 35

Polymerisation (A2)

Condensation polymers can be designed for specific properties — and, unlike addition polymers, they can be broken back down.

Picture itTwo very different design problems. Making a polymer strong is about what holds the chains to each other. Making it degradable is about what holds the chain to itself. Kevlar solves the first with hydrogen bonding; polylactic acid solves the second with hydrolysable ester links.

Polyesters and polyamides

Formed by condensation from a diol and dicarboxylic acid, or a diamine and dicarboxylic acid (or diacyl chloride, which is faster and irreversible). Terylene, nylon-6,6 and Kevlar are the standard examples, and you should be able to draw the repeat unit and identify the monomers from a chain.

Properties from intermolecular forces

Kevlar's exceptional tensile strength comes from extensive hydrogen bonding between aligned, rigid aromatic chains. Nylon has hydrogen bonding too, but its flexible aliphatic chains give a fibre rather than an armour material.

Degradability

Ester and amide links can be hydrolysed, so condensation polymers can be broken down chemically or biologically. Addition polymers have an unreactive C–C backbone and persist. Polylactic acid, made from renewable lactic acid, is designed to hydrolyse in the environment.

Recycling and disposal

Mechanical recycling requires sorting and downgrades quality. Chemical recycling depolymerises back to monomers and is better in principle but more expensive. Incineration recovers energy but requires gas treatment. Each is a trade-off, not a solution.

Designing for purpose

Chain length, branching, cross-linking, hydrogen bonding capability and rigidity of the monomer all affect strength, flexibility and melting behaviour — which means polymer properties are engineered rather than discovered.

The bit that catches people out'Biodegradable' does not mean 'degrades anywhere'. Polylactic acid needs industrial composting conditions — sustained high temperature and moisture — to break down in a useful time. In a landfill it can persist for years, which is a genuine and under-appreciated limitation.

The grown-up words

What it means	What it is called	Note
Polymer with ester links	polyester	Diol plus dicarboxylic acid
Polymer with amide links	polyamide	Diamine plus dicarboxylic acid
Aromatic polyamide of exceptional strength	Kevlar	Extensive hydrogen bonding
Breaking a polymer with water	hydrolysis	Only condensation polymers
Breaking a polymer back to monomers	chemical recycling	Better quality, higher cost
Polymer designed to break down naturally	biodegradable polymer	e.g. polylactic acid

Check you have got it

Why is Kevlar so strong for its mass?

Its rigid aromatic chains align closely and form extensive hydrogen bonds between adjacent chains, so a large force is needed to move chains past one another.

Give one reason chemical recycling is preferred to mechanical recycling where cost allows.

It breaks the polymer back to pure monomers, so the recycled material is as good as new, whereas mechanical recycling degrades chain length and quality with each cycle.

Edvia Free Resources · Chemistry 9701 · Topic 35 — free to copy and share

Topic 36

Organic synthesis (A2)

At A2 you plan multi-step routes and must justify them — including why you did not take the obvious one.

Picture itA chess problem rather than a maze. There is usually more than one legal route, and the best one accounts for yield, selectivity, cost and safety, not just the number of moves. Being asked to justify a route is being asked to think like a chemist rather than a lookup table.

The full reaction map

AS routes plus the A2 additions: arenes (nitration, halogenation, Friedel–Crafts), phenols, acyl chlorides, amines from nitriles and nitro compounds, diazotisation and coupling. Draw the map once yourself — the act of drawing it is what fixes it.

Chain length changes

Increase by one: halogenoalkane + KCN, or carbonyl + HCN, both giving nitriles. Increase by more: Friedel–Crafts. Decrease: rarely required at this level. Comparing the carbon counts of start and target is the first thing to do in any synthesis question.

Protecting groups and selectivity

When a molecule has two reactive groups, the reagent may attack the wrong one. Sometimes the order of steps solves this; sometimes a group must be temporarily protected. Recognising the problem is worth as much as solving it.

Yield across multiple steps

Overall yield is the product of the individual yields, so three steps at 70% give only 34% overall. This is why a shorter route with modest yields can beat a longer one with good yields — and why chemists care so much about step count.

Practical techniques and safety

Reflux, distillation, solvent extraction, recrystallisation, drying agents, thin-layer chromatography to monitor progress, melting point and spectroscopy to confirm the product. Hazards must be identified specifically — 'toxic' is not a risk assessment.

The bit that catches people outOverall yield multiplies, so it collapses fast. Four steps at 80% each give 41% overall, not 80%. When a question asks you to compare two routes, doing that arithmetic is usually the decisive argument.

The grown-up words

What it means	What it is called	Note
Planning backwards from the target	retrosynthesis	Standard planning method
Temporarily blocking a reactive group	protecting group	Removed later
Product of the individual step yields	overall yield	Falls quickly with steps
Reacting at one site in preference to another	selectivity	Often set by conditions
Monitoring a reaction's progress	thin-layer chromatography	Shows when it is complete
Confirming identity and purity	spectroscopy and melting point	Sharp mp means pure

Check you have got it

A synthesis has four steps with yields 90%, 80%, 75% and 60%. What is the overall yield?

0.90 × 0.80 × 0.75 × 0.60 = 0.324, so about 32%.

How would you convert benzene into phenylamine?

Nitrate with concentrated HNO₃ and H₂SO₄ at about 55 °C to give nitrobenzene, then reduce with tin and concentrated HCl, followed by NaOH to liberate the free amine.

Edvia Free Resources · Chemistry 9701 · Topic 36 — free to copy and share

Topic 37

Analytical techniques (A2)

Proton NMR is the most informative single technique in organic chemistry, because it maps the hydrogen atoms and their neighbours.

Picture itA group photograph where you can tell how many people are in each family, and how many people each family is standing next to. The chemical shift says what kind of environment; the integration says how many hydrogens; the splitting says how many hydrogens are on the adjacent carbon. Three pieces of information from one spectrum.

Reading a proton NMR spectrum

Number of signals = number of different hydrogen environments. Chemical shift (from the data sheet) indicates the type of environment. Integration gives the ratio of hydrogens in each environment. Splitting gives the number of hydrogens on adjacent carbons.

The n+1 rule

A signal split into n+1 peaks means n hydrogens on the neighbouring carbon. So a triplet means two neighbours, a quartet means three. The classic ethyl group pattern — a triplet and a quartet in a 3:2 ratio — is worth recognising instantly.

TMS and solvents

Tetramethylsilane is the reference at δ = 0: it is inert, volatile, non-toxic and gives a single sharp peak upfield of almost everything else. Deuterated solvents such as CDCl₃ are used because they contain no ordinary hydrogen to give signals.

D₂O exchange identifies OH and NH

Adding D₂O exchanges the labile hydrogens of OH and NH groups for deuterium, so those signals disappear. A peak that vanishes on shaking with D₂O is an OH or NH proton — a simple and decisive test.

Combining techniques

Mass spectrometry gives the molecular formula, infrared identifies the functional groups, carbon-13 NMR counts the carbon environments and proton NMR gives the hydrogen arrangement. Structure determination questions expect you to use all four in sequence and state what each contributes.

The bit that catches people outIntegration gives the ratio of hydrogens, not the absolute number. An integration of 3:2 could be C₅H₁₀ or C₁₀H₂₀ — you need the molecular formula from mass spectrometry to fix the actual count.

The grown-up words

What it means	What it is called	Note
Number of different hydrogen environments	number of signals	In proton NMR
Position of a signal relative to TMS	chemical shift	Indicates the environment
Relative area under each signal	integration	Gives the hydrogen ratio
Splitting into n+1 peaks	spin-spin coupling	n hydrogens on the neighbour
Reference compound at delta = 0	TMS	Inert, volatile, single peak
Shaking to remove OH and NH signals	D2O exchange	Labile protons swap for deuterium

Check you have got it

A signal appears as a quartet. How many hydrogens are on the adjacent carbon?

Three. By the n+1 rule, a quartet means n + 1 = 4, so n = 3 — typical of a CH₂ next to a CH₃.

A peak at δ 2.1 disappears when the sample is shaken with D₂O. What does this tell you?

It is an OH or NH proton — these are labile and exchange with deuterium, so the signal vanishes.

Edvia Free Resources · Chemistry 9701 · Topic 37 — free to copy and share

Chemistry 9701 — chapter handouts

Topics

Atomic structure

Subatomic particles and isotopes

Subshells fill in a specific order

Ionisation energy and what it reveals

The trends across and down

Orbital shapes

The grown-up words

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Atoms, molecules and stoichiometry

The three core equations

Empirical and molecular formulae

Limiting reagent and yield

Titration calculations

Balancing, including ionic and half equations

The grown-up words

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Chemical bonding

The three bond types

VSEPR predicts shape

Electronegativity and polarity

Intermolecular forces, weakest to strongest

Hydrogen bonding explains water's anomalies

The grown-up words

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States of matter

The ideal gas equation and where it fails

Simple molecular lattices

Giant covalent lattices

Ionic and metallic lattices

Predicting properties from structure

The grown-up words

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Chemical energetics

Enthalpy changes and their definitions

Hess's law and cycles

Bond energies give approximate answers

Calorimetry

Born–Haber cycles

The grown-up words

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Electrochemistry (AS)

Oxidation numbers and the rules

Half equations and combining them

Electrolysis

The Faraday relationships

Naming and predicting products

The grown-up words

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Equilibria (AS)

Dynamic equilibrium

Le Chatelier's principle

The equilibrium constant

Industrial compromise

Acids, bases and pH

The grown-up words

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Reaction kinetics (AS)

Collision theory

Factors affecting rate

The Boltzmann distribution

Catalysts

Following a reaction experimentally

The grown-up words

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Chemical periodicity

Atomic radius and ionisation energy

Melting points across period 3

Oxides across period 3

Chlorides across period 3

Group 2 and group 17 trends

The grown-up words

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Group 2 (AS)

Reactions with oxygen, water and acids

Reactivity increases down the group

Hydroxide solubility increases down the group

Thermal decomposition of carbonates and nitrates

Uses that follow from the chemistry