OCR A level Chemistry A (H032, H432) revision notes – full course
2.1 Atoms and Reactions
2.1.1 Atomic structure and isotopes
Isotopes are atoms of the same element with different numbers of neutrons and different masses. Atoms comprise of
protons, neutrons and electrons.
Particle Relative charge Relative mass Relative isotopic mass is the mass of an isotope compared to
Proton 1+ 1 1/12 of the mass of an atom of carbon-12. Relative atomic mass
Neutron 0 1 is the weighted, mean mass of an atom of an element compared
Electron 1- 1/1836 to 1/12 of the mass of carbon-12.
Mass spectrometry can be used to determine the relative isotopic masses and abundances of an isotope.
1. A sample is vaporised and ionised to form positive ions
2. The ions are accelerated. Heaver ions are more difficult to deflect than lighter ions and move more slowly, so the ions
are separated.
3. The ions are detected on a mass spectrum as a mass-to-charge ratio m/z. The greater the abundance, the larger the
signal.
2.1.2 Compounds, formulae and equations
The formulae of ionic compounds can be achieved by prediction of ionic charges from the position of elements in the
periodic table. There are also some common ions to remember: NO3-, CO32-, SO42-, OH-, NH4+, Zn2+ and Ag+. State symbols
are used to indicate physical state.
2.1.3 Amount of substance
Definitions:
1. Amount of substance: the quantity whose unit is the mole and which is used as a means of counting any species such
as atoms, ions and molecules.
2. Mole: the amount of any substance containing as many elementary particles as there are carbon atoms in exactly 12g
of the carbon-12 isotope – 6.02×1023 particles. It is abbreviated to mol and this is also its unit.
3. Avogadro constant, NA: the number of atoms per mole of the carbon-12 isotope. Unit: mol-1.
4. Molar mass, M: the mass of a mole of substance, in units of gmol-1.
5. Relative molecular mass, Mr: the weighted, mean mass of a molecule of a compound compared to 1/12 of the mass of
an atom of carbon-12. Unit: AMU (atomic mass unit).
6. Molar gas volume, Vm: gas volume per mole (24dm3 at RTP). Unit: dm3mol-1.
Formulae:
1. Empirical formula: the simplest whole number ratio of atoms of each element present in a compound.
2. Molecular formula: the number and type of atoms of each element in a molecule.
The empirical formula can be calculated from composition by mass through calculating the molar 𝑀𝑎𝑠𝑠
ratio. The molecular formula can be found from the empirical formula by multiplying each 𝑀𝑜𝑙𝑒𝑠 =
𝑅𝐹𝑀
component of the formula by relative molecular mass/empirical formula mass.
Hydrated salts:
1. Hydrated: a crystalline compound containing water molecules.
2. Anhydrous: a crystalline compound containing no water molecules.
3. Water of crystallisation: water molecules that are bonded into the crystalline structure of a hydrated compound. Can
be driven off by heating the crystalline compound.
The formula of a hydrated salt can be calculated from a given percentage composition, mass composition or experimental
result (i.e. where the mass of a compound before and after the water is driven off is recorded). This is done in a similar
way to the calculation of empirical formulae: the number of moles of the anhydrous salt is calculated, the number of moles
of water is calculated, and, finally, the smallest whole-number ratio between the two is found. Experimentally finding the
formula of hydrated salts is not very accurate because we assume that: 1) all the water has been lost, and 2) no further
decomposition has occurred (some salts decompose upon heating).
Calculations involving gas volume: 1cm3 = 1ml, 1m3 = 1,000dm3 = 1,000,000cm3
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,Use the equation:
𝑀𝑜𝑙𝑒𝑠 = 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 × 𝑣𝑜𝑙𝑢𝑚𝑒 or 𝑛 = 𝑐 × 𝑉
Standard solutions have a known concentration. They are prepared by dissolving an exact mass of solute into a solvent and
making up the solution to an exact volume.
Molar gas volume: At RTP (about 20°C and 101kPA, 1 atm), the molar gas volume = 24.0dm3mol-1.
Ideal gas equation: the ideal gas equation allows you to calculate gas volume etc. when the gas is not at RTP. You must
make some assumptions for the molecules making up an ideal gas:
▪ Random motion
▪ Elastic collisions
▪ Negligible size
▪ No intermolecular forces.
𝑝𝑉 = 𝑛𝑅𝑇
Where:
1. p is pressure (Pa)
2. V is volume (m3)
3. n is the amount of gas molecules (mol)
4. R is the ideal gas constant (8.31 Jmol-1K-1)
5. T is the temperature (K)
Stoichiometric relationships in calculations: the stoichiometry of a reaction is the ratio of amount, in moles, of each
substance in a chemical equation (essentially the ratio of the balancing numbers). This can be used to find the quantities of
the reactants required to prepare a quantity of product or the quantity of products that should be formed from a certain
quantity of reactants.
Percentage yield: the theoretical yield may not always be achieved because the reaction may not have gone to completion,
other reactions (side reactions) may have taken place and the purification of the product may result in the loss of some
product. The actual yield is usually lower than the theoretical yield. The conversion of starting materials into a desired
product is expressed by the percentage yield.
𝑎𝑐𝑡𝑢𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑦𝑖𝑒𝑙𝑑 = × 100%
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
Limiting reagents: in a reaction, the reactant that is not in excess will be completely used up first and stop the reaction. It is
called the limiting reagent. You can work out which reactant is in excess by calculating the amount of moles of each
reactant and comparing with the equation. Any subsequent calculation must be based on the limiting reagent.
Atom economy: the atom economy of a chemical reaction is a measure of how well the atoms have been utilised. Reactions
with high atom economies produce a large proportion of desired products with few waste products. Atom economy is
based solely on the balanced chemical equation of a reaction and assumes a 100% yield.
𝑠𝑢𝑚 𝑜𝑓 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠𝑒𝑠 𝑜𝑓 𝑑𝑒𝑠𝑖𝑟𝑒𝑑 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠
𝐴𝑡𝑜𝑚 𝑒𝑐𝑜𝑛𝑜𝑚𝑦 = × 100%
𝑠𝑢𝑚 𝑜𝑓 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠𝑒𝑠 𝑜𝑓 𝑎𝑙𝑙 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠
There are benefits for sustainability associated with developing processes with a high atom economy:
1. An awareness if atom economy allows manufacturers to make the best use of dwindling, finite resources and
minimising (harmful) waste which must be processed and disposed of.
2. Improving atom economy makes industrial processes more efficient, preserves raw materials and reduces waste.
3. Note that efficiency also depends on percentage yield, not just atom economy.
Techniques and procedures required during experiments requiring the measurement of mass, volumes of solutions or gas
volumes:
1. Measurement of mass: use a balance and avoid 0 errors. During the experiment to calculate the empirical formula of
MgO, ensure that the lid is either off or on consistently throughout the weighing processes.
2. Volumes of solutions: go down to eye level to avoid parallax. A syringe could be used.
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,3. Gas volumes: use a gas syringe, use downward delivery, use upward delivery or collect over water. If collecting over
water, note that any gas left in the delivery tubing is unimportant because the same volume of air will have been
originally displaced from the tubing into the collecting test tube at the start of the reaction.
2.1.4 Acids
Acids, bases, alkalis and neutralisation: acids release H+ ions in aqueous solution. Common examples include HCl, H2SO4,
HNO3 and CH3COOH. Alkalis release OH- ions in aqueous solutions. Common examples include NaOH, KOH and NH 3.
1. Strong acids release all their hydrogen atoms into solution as H+ ions; they completely dissociate in aqueous solution.
HCl H+ + Cl-
2. Weak acids only release a small proportion of the available hydrogen atoms as H+ ions in aqueous solution; they only
partially dissociate in aqueous solution. CH3COOH ⇌ H+ + CH3COO-
Note that concentration refers to the number of H+ ions per unit volume.
Neutralisation is the reaction of H+ and OH- ions to form H2O.
H+ + OH- H2O
It can also be described as the reaction of acids with bases, including carbonates, metal oxides and alkalis (water-soluble
bases) to form salts.
CuO + H2SO4 CuSO4 + H2O
NaOH + HCl NaCl + H2O
ZnCO3 + 2HCl ZnCl2 + CO2 + H2O
In these reactions, the hydrogen in the acid is replaced by a metal or ammonium ion in the salt.
Acid-base titrations
A titration is a technique used to find the concentration of a solution, identify an unknown chemical or find the purity of a
substance. It involves a number of procedures:
1. Preparing a standard solution: this is a solution of a known concentration. A certain mass of solid is dissolved in a
beaker using less distilled water than needed to fill the volumetric flask to the mark, then the solution is transferred to
the volumetric flask. The last traces if the solution are rinsed into the flask with the distilled water. The flask is
carefully filled to the graduation line by addition of distilled water a drop at a time until the bottom of the meniscus
lines up exactly with the mark. Then the volumetric flask is slowly inverted several times to mix the solution
thoroughly.
2. Carrying out the titration:
▪ Add a measured volume of one solution to a conical flask using a pipette (the aliquot).
▪ Add the other solution to the burette, and record the original burette reading.
▪ Add a few drops of indicator to the solution in the conical flask.
▪ Run the solution in the burette into the conical flask, swirling the conical flask throughout to mix the two
solutions. Eventually the indicator changes colour at the end point. This yields the volume of on solution required
to exactly react with volume of the second solution. Record the final burette reading and calculate the titre.
▪ Repeat accurately, adding the solution dropwise as the end point is approached. Repeat this until concordant
results (agree to within 0.10cm3) are achieved. This yields the mean titre.
2.1.5 Redox
Oxidation number:
The oxidation number is based upon a set of rules that apply to atoms; it can be thought of as the number of electrons
involved in the bonding to a different element.
1. Elements: the oxidation number is always 0 for elements.
2. Ions in compounds: equal to the modulus of the ionic charge with the sign in front of the number.
3. Combined atoms: sum of oxidation numbers = total charge
4. Special cases: the oxidation number of H in hydrides is -1; the oxidation number of O in peroxides is -1; the oxidation
number of O bonded to F is +2.
Roman numerals are used in the names of compounds of elements that form ions of different charges. The roman numeral
shows the oxidation state of the element (without a sign).
Chlorate(I) ClO-
Chlorate(III) ClO2-
3
, Nitrate(III) NO2-
Nitrate(V) NO3-
Redox reactions
Oxidation Reduction
Loss of electrons Gain of electrons
Addition of oxygen Removal of oxygen
Increase in oxidation number Decrease in oxidation number
Metals take part in redox reactions with acids to form salts and hydrogen gas.
Reaction of zinc with dilute hydrochloric acid
Zn + 2HCl ZnCl2 + H2
0 +2 Oxidation
+1 0 Reduction
Reaction of aluminium with dilute sulfuric acid
2Al + 3H2SO4 Al2(SO4)3 + 6H2
0 +3 Oxidation
+1 0 Reduction
This can be used to make predictions in terms of oxidation numbers and electron loss/gain in the context of unfamiliar
redox reactions: H+ has an oxidation number of +1, elements have oxidation numbers of 0 and there is no overall change
in oxidation number.
2.2 Electrons, bonding and structure
2.2.1 Electron structure
Shells, energy levels, or principal quantum numbers are made up of atomic orbitals – regions around the nucleus that can
hold up to two electrons with opposite spins - which in turn group to form sub-shells.
Types of orbital:
1. S-orbitals: spherical. The greater the shell number n, the greater the radius of its s-orbital. Sub-shells contain one s-
orbital.
2. P-orbitals: electron cloud is shape of a dumb-bell. Sub-shells contain three p-orbitals.
3. D-orbitals: more complex. Sub-shells contain five d-orbitals.
4. F-orbitals: also more complex. Sub-shells contain seven f-orbitals.
Filling of orbitals:
1. Orbitals fill in order of increasing energy. Within each shell, the new sub-shell added has a higher energy.
2. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f
3. The highest energy level in the third shell overlaps with the lowest energy level in the fourth shell.
4. Each orbital can hold up to two electrons. The two electrons in an orbital must have opposite spins to help counteract
the repulsion between the negative charges of the two electrons.
5. Within a sub shell, one electron occupies each orbital before pairing starts, preventing repulsion between paired
electrons until there is no further orbital available in the same energy level. At this point, electron pairing starts.
Electron configuration:
1. Express as above, using indices to show the number of electrons occupying each sub-shell.
2. Electron configurations can be expressed more simply in terms of the previous noble has in the periodic table plus the
outer electron sub shells. E.g. Lithium=1s22s1 = [He]2s1 in shorthand.
3. Electron configuration of ions: these are expressed in the same way as atoms – the electrons lost are omitted or the
electrons gained are added. For d-block elements, the 4s sub-shell empties before the 3d sub-shell (despite also filling
first).
4. Anomalies: Chromium atoms are given as [Ar]3d54s1. Copper atoms are given as [Ar]3d104s1 but copper ions are given
as [Ar]3d9.
Electron blocks:
The periodic table can be divided into blocks corresponding to the highest energy sub-shell.
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