AS Level Chemistry equations and key facts

Every chapter of AS Level Chemistry on one page: the 339 equations, definitions and facts to remember, in syllabus order. Use it for a last look before a test, then check yourself.

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Atomic structure

Particles in the atom and atomic radius

Proton: relative mass 1, relative charge +1. Neutron: relative mass 1, charge 0. Electron: relative mass 1/1840, charge −1.
Number of neutrons = nucleon number − proton number.
Electrons in an ion = proton number − charge on the ion (so a 3+ ion has 3 fewer electrons, a 2− ion has 2 more).
In an electric field: protons bend towards the negative plate, electrons bend towards the positive plate and much more (far smaller mass), neutrons go straight on.
Atomic radius decreases across a period (more protons, same shell, similar shielding) and increases down a group (more shells, more shielding).
A positive ion is smaller than its atom; a negative ion is larger than its atom.
For ions with the same number of electrons, the one with more protons is smaller.

Isotopes

Isotopes: same number of protons, different numbers of neutrons.
In ^x_yA: x = nucleon number (top), y = proton number (bottom).
Number of neutrons = x − y.
Same chemical properties, because the number and arrangement of electrons are the same.
Different physical properties (mass and density), because the atoms have different masses.
Atoms with the same nucleon number but different proton numbers are different elements, not isotopes.

Electrons, energy levels and atomic orbitals

s sub-shell: 1 orbital, 2 electrons. p: 3 orbitals, 6 electrons. d: 5 orbitals, 10 electrons.
Maximum electrons in a shell = 2n²: 2, 8, 18 for n = 1, 2, 3.
Filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p.
Chromium is [Ar] 3d⁵ 4s¹ and copper is [Ar] 3d¹⁰ 4s¹.
When atoms from scandium to zinc form positive ions, the 4s electrons are lost before the 3d electrons.
Orbitals of equal energy are each filled singly, with parallel spins, before any pairing.
Free radical: a species with one or more unpaired electrons, e.g. Cl•.

Ionisation energy

First ionisation energy: X(g) → X⁺(g) + e⁻. Second: X⁺(g) → X²⁺(g) + e⁻. Units: kJ mol⁻¹.
Across a period it generally increases: nuclear charge rises, same shell, similar shielding.
Down a group it decreases: the outer electron is further away and more shielded, which outweighs the larger nuclear charge.
Drop from Group 2 to Group 13 (Mg to Al): the electron removed is in a p sub-shell, higher in energy and more shielded.
Drop from Group 15 to Group 16 (P to S): spin-pair repulsion in a p orbital makes one electron easier to remove.
A big jump after the n^th ionisation energy means n outer-shell electrons.

Atoms, molecules and stoichiometry

Relative masses of atoms and molecules

Unified atomic mass unit (u) = one twelfth of the mass of a carbon-12 atom.
Relative atomic mass, A_r: the weighted mean mass of an atom of an element relative to one twelfth of the mass of an atom of carbon-12.
Relative isotopic mass: the mass of an atom of one isotope relative to one twelfth of the mass of an atom of carbon-12.
Relative molecular mass, M_r: the weighted mean mass of a molecule relative to one twelfth of the mass of an atom of carbon-12.
Relative formula mass: the same idea for one formula unit of an ionic compound.
M_r = sum of the A_r values of all the atoms in the formula. Multiply everything inside a bracket by the number outside.

The mole and the Avogadro constant

One mole contains 6.02 × 10²³ of the specified particles (L = 6.02 × 10²³ mol⁻¹).
amount in moles = mass ÷ molar mass (n = m/M), with mass in g and molar mass in g mol⁻¹.
number of particles = amount in moles × L.
number of atoms = moles of molecules × atoms in one molecule × L.
mass of one atom or molecule in grams = molar mass ÷ L.

Formulas

Ions to recall: NO₃⁻, CO₃²⁻, SO₄²⁻, OH⁻, NH₄⁺, Zn²⁺, Ag⁺, HCO₃⁻, PO₄³⁻.
Group 1 ions are 1+, Group 2 are 2+, Group 13 are 3+; Group 15 are 3−, Group 16 are 2−, Group 17 are 1−.
Empirical formula: divide each mass or percentage by A_r, then divide every answer by the smallest.
Molecular formula = empirical formula × (M_r ÷ empirical formula mass).
Balanced equation: the same number of each kind of atom, and the same total charge, on both sides.
State symbols: (s) solid, (l) liquid, (g) gas, (aq) dissolved in water.
For a hydrated salt, x = moles of water ÷ moles of anhydrous salt.

Reacting masses and volumes (of solutions and gases)

amount in moles = mass ÷ molar mass (n = m/M).
amount in moles = concentration × volume (n = cV), with volume in dm³; cm³ ÷ 1000 = dm³.
volume of gas = amount in moles × 24.0 dm³ at room conditions.
Gases at the same temperature and pressure react in the same volume ratio as the mole ratio in the equation.
Limiting reagent: divide the moles of each reactant by its number in the equation; the smallest result runs out first.
percentage yield = (actual yield ÷ theoretical yield) × 100.

Chemical bonding

Electronegativity and bonding

Electronegativity: the power of an atom to attract electrons to itself.
Increases across a period: nuclear charge increases, atomic radius decreases, shielding stays about the same.
Decreases down a group: the bonding electrons are further from the nucleus and more shielded.
Fluorine has the highest Pauling value (4.0).
Large difference in electronegativity (roughly 1.7 or more): ionic bond.
Smaller difference: covalent bond; the bigger the difference, the more polar the bond.

Ionic bonding

Ionic bonding: the electrostatic attraction between oppositely charged ions.
Sodium chloride: Na⁺ and Cl⁻, ratio 1 : 1. Each ion is surrounded by six of the other kind.
Magnesium oxide: Mg²⁺ and O²⁻, ratio 1 : 1. Two electrons are transferred from each Mg atom.
Calcium fluoride: Ca²⁺ and F⁻, ratio 1 : 2, formula CaF₂.
Total positive charge = total negative charge in the formula.
Higher charge and smaller ionic radius give stronger attraction and a higher melting point.

Metallic bonding

Metallic bonding: the electrostatic attraction between positive metal ions and delocalised electrons.
Number of delocalised electrons per atom = number of outer-shell electrons (Na 1, Mg 2, Al 3).
The bonding is stronger when the ion has a higher charge, there are more delocalised electrons and the ion is smaller.
Metals conduct electricity because the delocalised electrons are free to move; the ions do not move.
Metals are malleable because layers of ions slide over each other without the bonding breaking.
High melting and boiling points: much energy is needed to overcome the strong attraction.

Covalent bonding and coordinate (dative covalent) bonding

Single bond = 1 σ. Double bond = 1 σ + 1 π. Triple bond = 1 σ + 2 π.
Coordinate bond: a shared pair in which both electrons come from one atom, e.g. NH₃ + H⁺ → NH₄⁺.
Al₂Cl₆ has two coordinate bonds, each from a chlorine lone pair to an aluminium atom.
Hybridisation: sp³ in methane and ethane (109.5°), sp² in ethene (120°), sp in ethyne and CO₂ (180°).
Bond energy: the energy needed to break one mole of a particular covalent bond in the gaseous state.
Bond length: the distance between the nuclei of the two bonded atoms.
More shared pairs give a shorter and stronger bond: C≡C is shorter and stronger than C=C, which is shorter and stronger than C–C.

Shapes of molecules

2 bonding regions, no lone pairs: linear, 180° (CO₂).
3 bonding pairs, no lone pairs: trigonal planar, 120° (BF₃).
4 bonding pairs: tetrahedral, 109.5° (CH₄).
3 bonding pairs + 1 lone pair: pyramidal, 107° (NH₃). 2 bonding pairs + 2 lone pairs: non-linear, 104.5° (H₂O).
5 bonding pairs: trigonal bipyramidal, 120° and 90° (PF₅). 6 bonding pairs: octahedral, 90° (SF₆).
Repulsion: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.
For an ion, add one electron for each negative charge and remove one for each positive charge before counting pairs.

Intermolecular forces, electronegativity and bond properties

Electronegativity increases across a period and decreases down a group; fluorine is the most electronegative element.
Symmetrical molecules such as CO₂, BF₃, CH₄ and CCl₄ have polar bonds but no overall dipole, because the bond dipoles cancel.
id-id forces act between all molecules; they are stronger when the molecules have more electrons.
pd-pd forces are an extra attraction between molecules with a permanent dipole, e.g. HCl and propanone.
Hydrogen bonding needs an H atom bonded to N or O, and a lone pair on N or O, e.g. H₂O, NH₃ and alcohols.
One water molecule can form up to 4 hydrogen bonds; this gives water its high melting point, boiling point and surface tension.
Ice is less dense than water because hydrogen bonds hold the molecules in an open lattice, further apart than in the liquid.

Dot-and-cross diagrams

Outer electrons to start with: H 1, B 3, C 4, N and P 5, O and S 6, halogens 7.
Single bond = 1 shared pair (2 electrons); double bond = 2 shared pairs (4 electrons); triple bond = 3 shared pairs (6 electrons).
Lone pairs are outer electrons not used in bonding: N in NH₃ has 1 lone pair, O in H₂O has 2.
Coordinate bond: both electrons come from one atom, e.g. N to H⁺ in NH₄⁺, and Cl to Al in Al₂Cl₆.
Ionic: draw each ion in square brackets with the charge outside; the negative ion shows the electrons it gained with the other atom's symbol.
Fewer than eight: B in BF₃ has 6. Expanded octet: S in SO₂ has 10, P in PCl₅ has 10, S in SF₆ has 12.
Check: total electrons drawn = sum of the outer electrons of all the atoms, minus one for each positive charge, plus one for each negative charge.

States of matter

The gaseous state: ideal and real gases and pV = nRT

pressure × volume = amount in moles × gas constant × temperature (pV = nRT)
Units: p in Pa, V in m³, T in K, R = 8.31 J K⁻¹ mol⁻¹.
Conversions: 1 kPa = 1000 Pa; 1 dm³ = 10⁻³ m³; 1 cm³ = 10⁻⁶ m³; T in K = temperature in °C + 273.
To find M_r: first n = pV ÷ (RT), then M_r = mass in g ÷ n.
Ideal gas: zero particle volume and no intermolecular forces of attraction.
Real gases are most ideal at low pressure and high temperature.
Gases with strong intermolecular forces or large molecules (e.g. NH₃) deviate more than small, non-polar ones (e.g. He).

Bonding and structure

Giant ionic (NaCl, MgO): high melting point; conducts when molten or in solution, not when solid; often soluble in water.
Simple molecular (I₂, C₆₀, ice): low melting point; does not conduct in any state.
Giant molecular (diamond, SiO₂): very high melting point; insoluble; does not conduct.
Diamond: each C bonded to 4 others, tetrahedral. SiO₂: each Si bonded to 4 O, each O bonded to 2 Si.
Graphite: each C bonded to 3 others in layers; the fourth electron is delocalised, so graphite conducts; weak forces between layers.
Giant metallic (Cu): positive ions in a sea of delocalised electrons; conducts as a solid and as a liquid; malleable.
MgO melts at a higher temperature than NaCl: its ions have charges of 2+ and 2−, so they attract more strongly.

Chemical energetics

Enthalpy change, ΔH

Standard conditions (shown by ^⦵): 298 K and 101 kPa.
ΔH_f: enthalpy change when 1 mol of a compound is formed from its elements in their standard states.
ΔH_c: enthalpy change when 1 mol of a substance is burned completely in oxygen.
ΔH_neut: enthalpy change when 1 mol of water is formed from an acid and an alkali.
ΔH = total energy of bonds broken − total energy of bonds formed.
heat energy = mass × specific heat capacity × temperature change (q = mcΔT); m is the mass of the water or solution in g, c = 4.18 J g⁻¹ K⁻¹.
ΔH = −mcΔT ÷ n, where n is the amount in moles that reacted; divide by 1000 to change J to kJ.

Hess's law

Hess's law: the enthalpy change is independent of the route, provided the initial and final conditions are the same.
From formation data: ΔH = sum of ΔH_f of products − sum of ΔH_f of reactants.
From combustion data: ΔH = sum of ΔH_c of reactants − sum of ΔH_c of products.
From bond energies: ΔH = bonds broken − bonds formed.
ΔH_f of an element in its standard state is zero.
Multiply each value by the number of moles in the balanced equation.
Reversing a reaction changes the sign of ΔH.

Electrochemistry

Redox processes: electron transfer and changes in oxidation number

Oxidation: loss of electrons, oxidation number increases. Reduction: gain of electrons, oxidation number decreases.
An uncombined element has oxidation number 0; a simple ion has an oxidation number equal to its charge.
The oxidation numbers add up to 0 in a compound and to the charge in an ion.
Usual values: Group 1 metals +1, Group 2 metals +2, F −1, H +1 (−1 in metal hydrides), O −2 (−1 in peroxides).
Oxidising agent: accepts electrons and is reduced. Reducing agent: donates electrons and is oxidised.
A Roman numeral shows the oxidation number: iron(III) is Fe³⁺, chlorate(I) is ClO⁻, manganate(VII) is MnO₄⁻.
To balance a redox equation: total increase in oxidation number = total decrease in oxidation number.

Equilibria

Chemical equilibria: reversible reactions, dynamic equilibrium

For aA + bB ⇌ cC + dD: K_c = [C]^c[D]^d ÷ ([A]^a[B]^b), with concentrations in mol dm⁻³ at equilibrium.
mole fraction = moles of one gas ÷ total moles of gas; partial pressure = mole fraction × total pressure.
K_p has the same form as K_c but uses partial pressures of the gases.
Raising the temperature moves the equilibrium in the endothermic direction; lowering it favours the exothermic direction.
Raising the pressure moves the equilibrium to the side with fewer moles of gas.
A catalyst does not change the position of equilibrium or the value of K; equilibrium is just reached faster.
Haber process: about 450 °C, about 200 atm, iron catalyst. Contact process: about 450 °C, 1–2 atm, V₂O₅ catalyst.

Brønsted–Lowry theory of acids and bases

Acid = proton donor; base = proton acceptor.
Strong acids: HCl, H₂SO₄, HNO₃. Weak acid: CH₃COOH. Strong bases: NaOH, KOH. Weak base: NH₃.
At the same concentration a strong acid has a lower pH, a higher conductivity and a faster reaction with a metal than a weak acid.
Neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l); acid + alkali → salt + water.
Vertical section of the pH curve: strong acid–strong alkali about pH 3 to 11; weak acid–strong alkali about 7 to 11; strong acid–weak alkali about 3 to 7.
An indicator is suitable if its colour-change range lies inside the vertical section of the curve.
Methyl orange (pH 3.1–4.4) suits a strong acid with a weak alkali; phenolphthalein (pH 8.3–10.0) suits a weak acid with a strong alkali; either suits strong with strong.

Reaction kinetics

Rate of reaction

rate = change in concentration ÷ time taken; unit mol dm⁻³ s⁻¹.
Rate can also be measured as a change in volume of gas or in mass per unit time, e.g. cm³ s⁻¹ or g s⁻¹.
Average rate = total change ÷ total time.
Rate at one moment = gradient of the tangent to the concentration–time graph at that time.
Effective collision: energy at least equal to the activation energy, and the correct orientation.
Higher concentration or higher gas pressure: more particles per unit volume, so a higher frequency of effective collisions.

Effect of temperature on reaction rates and the concept of activation energy

Activation energy, E_A: the minimum energy required for a collision to be effective.
Boltzmann curve axes: energy on the x-axis, number of molecules with that energy on the y-axis.
The curve starts at the origin, rises to a peak (the most probable energy) and never touches the energy axis again.
Area under the whole curve = total number of molecules; area to the right of E_A = molecules with enough energy to react.
Higher temperature: the peak is lower and further to the right; the total area stays the same.
Higher temperature: a greater proportion of molecules have energy ≥ E_A, so the rate increases.
Changing the temperature does not change E_A.

Homogeneous and heterogeneous catalysts

Catalyst: speeds up a reaction by providing an alternative route of lower activation energy; it is not used up.
Boltzmann curve: same curve, E_A marked further left, so a bigger area lies to the right of E_A.
Pathway diagram: lower peak for the catalysed route; ΔH is unchanged.
E_A of the forward and reverse reactions are lowered by the same amount.
A catalyst does not change the position of equilibrium or the yield; equilibrium is reached sooner.
Heterogeneous: different phase, e.g. solid V₂O₅ in the Contact process, solid iron in the Haber process.
Homogeneous: same phase, e.g. NO₂ gas catalysing the oxidation of SO₂ gas in the atmosphere.

The Periodic Table: chemical periodicity

Periodicity of physical properties of the elements in Period 3

Atomic radius decreases from Na to Cl: more protons, same number of shells, similar shielding.
Positive ions Na⁺, Mg²⁺, Al³⁺ have 2 shells and get smaller as nuclear charge rises; Al³⁺ is the smallest.
Negative ions P³⁻, S²⁻, Cl⁻ have 3 shells, so they are much larger than the positive ions; they also get smaller from P³⁻ to Cl⁻.
Melting point rises Na < Mg < Al: higher ion charge and more delocalised electrons give stronger metallic bonding.
Si has the highest melting point: giant molecular, many strong covalent bonds to break.
P₄, S₈, Cl₂ and Ar have low melting points (only id-id forces to overcome); order S₈ > P₄ > Cl₂ > Ar, following the number of electrons.
Conductivity increases Na < Mg < Al; Si is a semiconductor; P, S, Cl and Ar do not conduct.

Periodicity of chemical properties of the elements in Period 3

With oxygen: 4Na + O₂ → 2Na₂O; 2Mg + O₂ → 2MgO; 4Al + 3O₂ → 2Al₂O₃; P₄ + 5O₂ → P₄O₁₀; S + O₂ → SO₂
With chlorine: 2Na + Cl₂ → 2NaCl; Mg + Cl₂ → MgCl₂; 2Al + 3Cl₂ → Al₂Cl₆; Si + 2Cl₂ → SiCl₄; 2P + 5Cl₂ → 2PCl₅
With water: 2Na + 2H₂O → 2NaOH + H₂ (cold water, vigorous); Mg + H₂O(g) → MgO + H₂ (steam)
Oxides in water: Na₂O gives NaOH (pH 13–14); MgO reacts slightly (pH about 9); Al₂O₃ and SiO₂ do not react (pH 7); P₄O₁₀ + 6H₂O → 4H₃PO₄; SO₂ + H₂O → H₂SO₃; SO₃ + H₂O → H₂SO₄ (pH 1–3)
Chlorides in water: NaCl dissolves (pH 7); MgCl₂ dissolves (pH about 6.5); Al₂Cl₆ is hydrolysed (pH about 3); SiCl₄ + 2H₂O → SiO₂ + 4HCl; PCl₅ + 4H₂O → H₃PO₄ + 5HCl (pH 1–2)
Amphoteric: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O and Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄; Al(OH)₃ behaves in the same way
Highest oxidation number = number of outer-shell electrons: Na +1, Mg +2, Al +3, Si +4, P +5, S +6

Chemical periodicity of other elements

Same group = same number of outer-shell electrons = same formulas (for example all Group 2 oxides are MO, all Group 2 chlorides are MCl₂)
Highest oxidation number = number of outer-shell electrons (Group 14: +4, Group 16: +6)
Down a group: atomic radius increases and first ionisation energy decreases, so metals react more vigorously
Down Group 17: the elements are less volatile, darker in colour and weaker oxidising agents
Metal oxides are basic; non-metal oxides are acidic; oxides of elements near the metal/non-metal boundary (such as Al) are amphoteric
Ionic chlorides have high melting points and dissolve to give neutral solutions; covalent chlorides have low melting points and are hydrolysed to give acidic solutions with fumes of HCl
Down Group 2: hydroxides become more soluble, sulfates become less soluble, carbonates and nitrates become more stable to heat

Group 2

Similarities and trends in the properties of the Group 2 metals, magnesium to barium

With oxygen: 2M + O₂ → 2MO (white solid)
With water: M + 2H₂O → M(OH)₂ + H₂; magnesium is very slow in cold water, but with steam Mg + H₂O(g) → MgO + H₂
With acids: M + 2HCl → MCl₂ + H₂; M + H₂SO₄ → MSO₄ + H₂ (insoluble CaSO₄, SrSO₄ and BaSO₄ coat the metal and slow or stop the reaction)
Oxides with water: MO + H₂O → M(OH)₂; the solution is more alkaline down the group (Mg(OH)₂ pH about 9, Ca(OH)₂ pH about 12)
Oxides, hydroxides and carbonates neutralise acids: MO + 2HCl → MCl₂ + H₂O; M(OH)₂ + 2HCl → MCl₂ + 2H₂O; MCO₃ + 2HCl → MCl₂ + CO₂ + H₂O
Heating: MCO₃ → MO + CO₂; 2M(NO₃)₂ → 2MO + 4NO₂ + O₂ (brown fumes of NO₂); a higher temperature is needed down the group
Solubility down the group: hydroxides increase, sulfates decrease (BaSO₄ is insoluble)

Group 17

Physical properties of the Group 17 elements

Chlorine: yellow-green gas; bromine: red-brown liquid (orange-brown vapour); iodine: grey-black solid (purple vapour)
Volatility decreases down the group: boiling point Cl₂ < Br₂ < I₂
More electrons in the molecule = stronger id-id forces = higher boiling point
Bond energy decreases: Cl–Cl 242, Br–Br 193, I–I 151 kJ mol⁻¹
Larger atoms = longer bond = weaker attraction between the nuclei and the shared pair
Melting, boiling and subliming overcome forces between molecules; the X–X covalent bond is not broken

The chemical properties of the halogen elements and the hydrogen halides

Oxidising power: Cl₂ > Br₂ > I₂
Displacement: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂; Cl₂ + 2I⁻ → 2Cl⁻ + I₂; Br₂ + 2I⁻ → 2Br⁻ + I₂
Colours in water: bromine is orange-yellow, iodine is brown
With hydrogen: chlorine explodes in sunlight, bromine needs heating, iodine reacts slowly and reversibly (H₂ + I₂ ⇌ 2HI)
Bond energy: H–Cl 431, H–Br 366, H–I 299 kJ mol⁻¹
Thermal stability: HCl > HBr > HI; a hot rod decomposes HI (2HI → H₂ + I₂, purple fumes) but not HCl

Some reactions of the halide ions

Reducing power: Cl⁻ < Br⁻ < I⁻
Silver nitrate test: Ag⁺(aq) + X⁻(aq) → AgX(s)
AgCl: white, dissolves in dilute NH₃(aq); AgBr: cream, dissolves only in concentrated NH₃(aq); AgI: yellow, does not dissolve in either
Chloride with concentrated H₂SO₄: NaCl + H₂SO₄ → NaHSO₄ + HCl (steamy fumes, no redox)
Bromide: 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O (sulfur +6 to +4)
Iodide: 2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O; 6HI + H₂SO₄ → 3I₂ + S + 4H₂O; 8HI + H₂SO₄ → 4I₂ + H₂S + 4H₂O
Sulfur oxidation numbers: H₂SO₄ +6, SO₂ +4, S 0, H₂S −2

The reactions of chlorine

Disproportionation: one element is both oxidised and reduced in the same reaction
Cold dilute NaOH: Cl₂ + 2NaOH → NaCl + NaClO + H₂O (chlorine 0 to −1 and +1)
Hot concentrated NaOH: 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O (chlorine 0 to −1 and +5)
Water purification: Cl₂ + H₂O ⇌ HCl + HOCl (chlorine 0 to −1 and +1)
HOCl ⇌ H⁺ + ClO⁻; chlorine is +1 in both, so this step is not redox
NaClO is sodium chlorate(I); NaClO₃ is sodium chlorate(V)

Nitrogen and sulfur

N≡N bond energy is about 945 kJ mol⁻¹; N₂ is non-polar
NH₃ + H⁺ → NH₄⁺: the fourth N–H bond is a dative (coordinate) bond; the ion is tetrahedral with bond angles of 109.5°
Ammonium salt + alkali, warmed: NH₄⁺ + OH⁻ → NH₃ + H₂O
Formed in lightning (natural) and in hot engines and furnaces (man-made): N₂ + O₂ → 2NO, then 2NO + O₂ → 2NO₂
Catalytic converter: 2NO + 2CO → N₂ + 2CO₂
Acid rain: directly, 4NO₂ + 2H₂O + O₂ → 4HNO₃; as a catalyst, SO₂ + NO₂ → SO₃ + NO, then 2NO + O₂ → 2NO₂, and SO₃ + H₂O → H₂SO₄
NO and NO₂ + unburned hydrocarbons in sunlight → peroxyacetyl nitrate (PAN), a component of photochemical smog

An introduction to AS Level organic chemistry

Formulas, functional groups and the naming of organic compounds

Stems: meth- 1, eth- 2, prop- 3, but- 4, pent- 5, hex- 6 carbon atoms
General formulas: alkanes C_nH_2n+2; alkenes C_nH_2n; alcohols C_nH_2n+1OH; aldehydes and ketones C_nH_2nO; carboxylic acids and esters C_nH_2nO₂
Suffixes: -ene (C=C), -ol (OH), -al (CHO), -one (C=O in the chain), -oic acid (COOH), -nitrile (C≡N), -amine (NH₂); halogens are prefixes: chloro-, bromo-, iodo-
Number the longest chain from the end that gives the functional group the lowest number
The carbon atom of CHO, COOH and C≡N counts as carbon 1 of the chain
Esters: the alkyl part from the alcohol comes first, then the -oate part from the acid, so CH₃COOCH₃ is methyl ethanoate
Empirical formula = simplest whole-number ratio of the atoms in the molecular formula

Characteristic organic reactions

Homologous series: same functional group and general formula; each member differs from the next by CH₂
Saturated: only single carbon–carbon bonds; unsaturated: contains C=C or C≡C
Homolytic fission: each atom takes one electron, giving two free radicals (Cl₂ → 2Cl•); heterolytic fission: one atom takes both electrons, giving a positive and a negative ion
Free-radical substitution has three stages: initiation (radicals made), propagation (a radical is used and another is made), termination (two radicals join)
Nucleophile: electron-pair donor, e.g. OH⁻, CN⁻, NH₃; electrophile: electron-pair acceptor, e.g. H⁺, the H^δ+ of HBr
Hydrolysis: a bond is broken by reaction with water; condensation: two molecules join and a small molecule such as H₂O is lost
Mechanisms: alkanes, free-radical substitution; alkenes, electrophilic addition; halogenoalkanes, nucleophilic substitution; aldehydes and ketones, nucleophilic addition

Shapes of organic molecules; σ and π bonds

sp³ carbon: four σ bonds, tetrahedral, bond angle 109.5° (as in ethane)
sp² carbon: three σ bonds and one π bond, trigonal planar, bond angle 120° (as in ethene)
sp carbon: two σ bonds and two π bonds, linear, bond angle 180° (as in ethyne)
Single bond = 1 σ; double bond = 1 σ + 1 π; triple bond = 1 σ + 2 π
σ bond: end-on overlap of orbitals; π bond: sideways overlap of p orbitals
Ethene is planar: both carbon atoms and all four hydrogen atoms lie in the same plane

Isomerism: structural isomerism and stereoisomerism

Chain isomers: different carbon skeleton (butane and 2-methylpropane)
Positional isomers: same functional group in a different position (propan-1-ol and propan-2-ol)
Functional group isomers: different functional groups (propanal and propanone, both C₃H₆O)
Cis/trans isomerism needs a C=C bond with two different groups on each of its carbon atoms; cis = same side, trans = opposite sides
A ring also restricts rotation, so cyclic compounds can show cis/trans isomerism
Chiral centre: a carbon atom with four different groups attached; it gives two optical isomers (enantiomers)
A molecule can have more than one chiral centre

Hydrocarbons

Alkanes

Hydrogenation: alkene + H₂ → alkane, with a Pt or Ni catalyst and heat, e.g. CH₂=CH₂ + H₂ → CH₃CH₃
Cracking: heat with an Al₂O₃ catalyst; long alkane → shorter alkane + alkene
Complete combustion (excess oxygen) gives CO₂ and H₂O: CH₄ + 2O₂ → CO₂ + 2H₂O
Incomplete combustion (limited oxygen) gives toxic CO or carbon (soot): 2CH₄ + 3O₂ → 2CO + 4H₂O
Initiation: Cl₂ → 2Cl• (ultraviolet light, homolytic fission)
Propagation: CH₄ + Cl• → CH₃• + HCl, then CH₃• + Cl₂ → CH₃Cl + Cl•
Termination: two radicals join, e.g. CH₃• + Cl• → CH₃Cl or CH₃• + CH₃• → C₂H₆

Alkenes

Making alkenes: halogenoalkane + NaOH in ethanol, heat (elimination); alcohol + heated Al₂O₃ or concentrated H₂SO₄ (dehydration); cracking of a longer alkane
Additions: H₂ with Pt or Ni, heat → alkane; steam with H₃PO₄ catalyst → alcohol; HX(g) at room temperature → halogenoalkane; X₂ → dihalogenoalkane
Test for C=C: aqueous bromine changes from orange to colourless
Cold, dilute, acidified KMnO₄ → diol, e.g. ethene → ethane-1,2-diol
Hot, concentrated, acidified KMnO₄ breaks the C=C bond: a =CH₂ end gives CO₂, a =CHR end gives RCOOH, a =CR₂ end gives a ketone
Mechanism: the π bond attacks H^δ+ of H–Br (or the induced Br^δ+ of Br₂), a carbocation forms, then Br⁻ attacks the carbocation
Carbocation stability: tertiary > secondary > primary

Halogen compounds

Halogenoalkanes

Primary, secondary, tertiary: one, two or three alkyl groups on the carbon atom bonded to the halogen
NaOH(aq), heat → alcohol; KCN in ethanol, heat → nitrile (one more carbon); NH₃ in ethanol, heated under pressure → amine
NaOH in ethanol, heat → alkene (elimination of HX)
Bond strength: C–F > C–Cl > C–Br > C–I, so rate of substitution: iodo > bromo > chloro
With aqueous silver nitrate in ethanol: AgCl white, AgBr cream, AgI yellow precipitate
Primary: S_N2. Tertiary: S_N1. Secondary: a mixture of both
Making them: alkane + Cl₂ or Br₂ in ultraviolet light; alkene + X₂ or HX(g); alcohol + HX(g), PCl₅ at room temperature, PCl₃ and heat, or SOCl₂

Hydroxy compounds

Alcohols

Making alcohols: alkene + steam with H₃PO₄ catalyst; halogenoalkane + NaOH(aq), heat; aldehyde or ketone + NaBH₄ or LiAlH₄; carboxylic acid + LiAlH₄; hydrolysis of an ester
Primary alcohol + acidified K₂Cr₂O₇: distil → aldehyde; reflux with excess → carboxylic acid
Secondary alcohol → ketone; tertiary alcohol → no reaction, solution stays orange
With sodium: 2C₂H₅OH + 2Na → 2C₂H₅ONa + H₂
Dehydration: heated Al₂O₃ or concentrated H₂SO₄ → alkene + H₂O
To a halogenoalkane: HX(g), PCl₅, PCl₃ and heat, or SOCl₂. To an ester: carboxylic acid with concentrated H₂SO₄ catalyst
Alkaline I₂(aq) gives a yellow precipitate of tri-iodomethane, CHI₃, only if the alcohol has a CH₃CH(OH)– group

Carbonyl compounds

Aldehydes and ketones

Making them: primary alcohol + acidified K₂Cr₂O₇, distil → aldehyde; secondary alcohol + acidified K₂Cr₂O₇ → ketone
2,4-DNPH gives an orange precipitate with both aldehydes and ketones
Tollens' reagent: silver mirror with an aldehyde, no change with a ketone
Fehling's solution: blue solution gives a red-brown precipitate with an aldehyde, no change with a ketone
Reduction with NaBH₄ or LiAlH₄: aldehyde → primary alcohol; ketone → secondary alcohol
HCN with KCN catalyst, heat → hydroxynitrile, e.g. CH₃CHO + HCN → CH₃CH(OH)CN
Alkaline I₂(aq) gives a yellow precipitate of CHI₃ only with a CH₃CO– group (ethanal and methyl ketones)

Carboxylic acids and derivatives

Carboxylic acids

With a reactive metal: 2CH₃COOH + Mg → (CH₃COO)₂Mg + H₂
With an alkali: CH₃COOH + NaOH → CH₃COONa + H₂O
With a carbonate: 2CH₃COOH + Na₂CO₃ → 2CH₃COONa + H₂O + CO₂
With an alcohol and concentrated H₂SO₄ catalyst: acid + alcohol ⇌ ester + water
Reduction: LiAlH₄ → primary alcohol
Making them: primary alcohol or aldehyde + excess acidified K₂Cr₂O₇ or KMnO₄, reflux
Also made by hydrolysis of a nitrile or an ester with dilute acid, or with dilute alkali followed by acidification

Esters

Making an ester: RCOOH + R'OH ⇌ RCOOR' + H₂O, with concentrated H₂SO₄ catalyst and heat
Name: alkyl part from the alcohol, -oate part from the acid; in RCOOR' the acid part is written first
Acid hydrolysis (dilute acid, heat under reflux): CH₃COOC₂H₅ + H₂O ⇌ CH₃COOH + C₂H₅OH
Alkaline hydrolysis (NaOH(aq), heat under reflux): CH₃COOC₂H₅ + NaOH → CH₃COONa + C₂H₅OH
Alkaline hydrolysis is not reversible because the carboxylate ion does not react with the alcohol
1 mol of ester gives 1 mol of alcohol and 1 mol of acid or salt

Nitrogen compounds

Primary amines

Reagent and conditions: NH₃ in ethanol, heated under pressure in a sealed tube
Overall equation: CH₃CH₂Br + 2NH₃ → CH₃CH₂NH₂ + NH₄Br
First NH₃ is the nucleophile; second NH₃ is the base
The carbon chain keeps the same length, and –NH₂ takes the position of the halogen
Excess ammonia reduces further substitution
percentage yield = actual mass ÷ theoretical mass × 100

Nitriles and hydroxynitriles

Nitrile: CH₃CH₂Br + KCN → CH₃CH₂CN + KBr (KCN in ethanol, heat)
Hydroxynitrile from an aldehyde: CH₃CHO + HCN → CH₃CH(OH)CN, 2-hydroxypropanenitrile
Hydroxynitrile from a ketone: (CH₃)₂CO + HCN → (CH₃)₂C(OH)CN, 2-hydroxy-2-methylpropanenitrile
HCN is a weak acid, so KCN is added to supply the CN⁻ nucleophile
Acid hydrolysis: CH₃CN + 2H₂O + HCl → CH₃COOH + NH₄Cl
Alkaline hydrolysis: CH₃CN + NaOH + H₂O → CH₃COONa + NH₃, then acidify to get CH₃COOH

Polymerisation

Addition polymerisation

Ethene: n CH₂=CH₂ → –(CH₂–CH₂)_n–, poly(ethene)
Chloroethene, CH₂=CHCl, gives poly(chloroethene), PVC, with repeat unit –CH₂–CHCl–
Repeat unit from monomer: change C=C to C–C and draw one bond outwards from each of the two carbons
Monomer from polymer: take two neighbouring chain carbons with their side groups and put the double bond back between them
The polymer has the same empirical formula as its monomer
Burning: incomplete combustion gives toxic carbon monoxide; PVC also gives acidic hydrogen chloride gas

Organic synthesis

Common chain of changes: alkene → halogenoalkane → alcohol → aldehyde or ketone → carboxylic acid
To add one carbon atom: halogenoalkane + KCN in ethanol, heat → nitrile; then hydrolyse → carboxylic acid
NaOH in water gives substitution (alcohol); NaOH in ethanol gives elimination (alkene)
Primary alcohol: distil with acidified K₂Cr₂O₇ → aldehyde; reflux with excess → carboxylic acid
Tests: C=C decolourises aqueous bromine; –COOH gives CO₂ with Na₂CO₃; C=O gives an orange precipitate with 2,4-DNPH; –CHO gives a silver mirror with Tollens' reagent
Acidified K₂Cr₂O₇ turns green with primary and secondary alcohols and with aldehydes
By-products: the minor isomer when HX adds to an unsymmetrical alkene; further substitution products in free-radical substitution

Analytical techniques

Infrared spectroscopy

O–H in an alcohol: broad absorption at 3200–3600 cm⁻¹
O–H in a carboxylic acid: very broad absorption at 2500–3000 cm⁻¹
C=O (aldehyde, ketone, carboxylic acid, ester): strong, sharp absorption at 1680–1750 cm⁻¹
C≡N in a nitrile: sharp absorption near 2250 cm⁻¹
N–H in an amine: absorption at 3300–3500 cm⁻¹
C=C in an alkene: absorption near 1650 cm⁻¹
C–O (alcohol, ester): 1040–1300 cm⁻¹; C–H: 2850–2950 cm⁻¹, present in almost every organic compound

Mass spectrometry

relative atomic mass = sum of (isotope mass × percentage abundance) ÷ 100
Molecular ion peak M: the peak of highest m/e, ignoring the small isotope peaks; its m/e = M_r
number of carbon atoms, n = (100 × abundance of [M + 1]) ÷ (1.1 × abundance of M)
One chlorine atom: M and [M + 2] in the ratio 3 : 1 (³⁵Cl and ³⁷Cl)
One bromine atom: M and [M + 2] in the ratio 1 : 1 (⁷⁹Br and ⁸¹Br)
Two bromine atoms: M, [M + 2], [M + 4] in the ratio 1 : 2 : 1
Common fragments: CH₃⁺ 15, C₂H₅⁺ 29, CH₂OH⁺ 31, CH₃CO⁺ or C₃H₇⁺ 43