OCR A-Level Chemistry
Rushali Chawan 13BWD
2.1 Atomic structure and isotopes
• Atomic number = number of protons (also electrons for a natural atom) in that element and the IDENTITY of element.
• Mass number = number of protons + number of neutrons.
• Relative mass of electron = 1/1836 → relative mass of proton/neutron = 1
• There is lots of empty space around the nucleus – neutrons provide the glue that holds nucleus together, despite the electrostatic
repulsion between its positively charged protons.
• The number of electrons = the number of protons for a neutral atom.
• Isotope = atoms of the same element with different numbers of neutrons and different masses.
• Chemical reactions: different isotopes of the same element have the same number of electrons.
• The number of neutrons has no effect on reactions of an element = different isotopes of an element react in the same way.
• Ions = atoms that have lost or gained electrons – same number of protons, different number of electrons.
2.2 Relative mass
• Mass defect = the small amount of mass lost caused by the strong nuclear force holding protons and neutrons together.
• The mass of a carbon-12 isotope is defined as 12 atomic mass units (12u) – the standards mass for atomic mass is 1u, the mass of
1/12th of an atom of carbon-12 – 1u = ~ mass of a proton or neutron.
• Relative isotopic mass = the mass of an isotope compared with 1/12th mass of a carbon-12 atom.
• The mass of a carbon-12 isotope is defined as 12 atomic mass units (12u).
• Relative atomic mass = the weighted mean mass of an atom of an element compared with 1/12th the mass of a carbon-12 atom. It is
like the mass of an ‘average’ atom – appears on the Periodic Table.
• The weighted mean mass takes account of: (1) % abundance of each isotope + (2) relative isotopic mass of each isotope.
Determination of relative atomic mass:
A mass spectrometer is used to find the % abundance of the isotope in a sample.
1. A sample is placed in the mass spectrometer.
2. The sample is vaporised and ionised to from positive ions.
- The temperature is increased, or the pressure is decreased.
3. The ions are accelerated. Heavier ions are more slow and difficult to deflect than lighter ions, so the ions of each isotope are
separated.
- Amount of deflection depends of the mass + charge. Least mass/ most charge = more deflection.
1. The ions are detected on a mass spectrum as a mass-to-charge ratio (m/z). Each ion reaching the detectors adds to the signal, so the
greater the abundance, the larger the signal.
Example: Relative atomic mass of chlorine
The mass spectrometer can also record the m/z ratio for each isotope, so that accurate values of relative isotopic mass can be measured.
Example – Mass spectra of chlorine and bromide molecules:
• Molecular ion peak = the peak with the highest m/z ratio = relative molecular mass of the compound.
• Sometimes, there are small peaks at M + 1: increased mass due to the presence of one C-13 atom or H.
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,OCR A-Level Chemistry
Rushali Chawan 13BWD
2.3 Formulae and equations
• Binary compounds = contains two elements only, e.g. Na and O form NaO.
• Polyatomic ions = ions that contain atoms of more than one element bonded together.
3.1 Amount of substance and the mole
• One mole = the amount of substance than contains 6.02 x 1023 particles.
• Avogadro constant is 6.02 x 1023mol-1, the number of particles in each mole of carbon-12.
• 12g of carbon-12 contains 6.02 x 1023 atoms + 1 mole of magnesium atoms has a mass of 24.3g.
• Molar mass = mass per mole of a substance (g mol-1).
3.2 Determination of formulae
• Molecular formulae = the number of atoms of each element in a molecule
• Empirical formulae = the simplest whole number ratio of atoms of each element in a
compound.
• Relative molecular mass (Mr) = compares the mass of a molecule with the mass of an
atom of C-12.
• Relative formula mass = compares the mass of a formula unit with the mass of a C-12
atom (add up the relative atomic masses).
• Water of crystallisation = water molecules are part of their crystalline structure.
Accuracy of experimental formulae:
1. Assumption 1 – all of the water has been lost.
- Some water may be left inside.
- So, heat to a constant mass.
2. Assumption 2 – No further decomposition.
- Many salts decompose further when heated, but with no colour change.
3.3 Moles and volumes
• The cubic centimetre (cm3) or millilitre (ml): 1cm3 = 1ml
• The cubic decimetre (dm3) or litre (l): 1dm3 = 1000cm3 = 1000ml = 1l
• Moles and solutions: a 1 mol dm-3 solution contains 1 mol of solute dissolved in each 1 dm3 of solution.
• Avogadro’s Law = equal volumes of gases under the same conditions of temperature + pressure contain the same no. of molecules.
• Mass concentration will often be in g cm-3, e.g. Na2CO3 has a concentration of 0.250 mol dm-3.
• Moles and gas volumes: at the same temperature/pressure, equal volumes of different gases contain the same number of molecules.
• Molar volume: molar gas volume (Vm) = the volume per mole of gas molecules at a stated temperature and pressure (dm3 mol-1).
• At RTP, 1 mole of gas molecules has a volume of ~24.0 dm3 = 24,000 cm3 → at RTP, the molar gas volume = 24.0 dm3 mol-1.
The ideal gas equation: Assumptions:
[1 atm = 101kPa = 101,000 Pa]
1. Forces between molecules are negligible.
2. Gas molecules have a negligible size compared to size of container.
3. Random movement of molecules.
4. Elastic collisions.
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,OCR A-Level Chemistry
Rushali Chawan 13BWD
3.3 Moles and volumes
E.g finding a relative molecular mass of a volatile liquid, which is liquid at RT + has a BP below 100˚C so that it vaporises:
3.4 Reacting quantities
• To determine an unknown G2 metal, react with HCl +measure maximum volume of H2 gas produced.
Percentage yield:
• The reaction many not be complete/finished.
• Other (side) reactions may have taken place alongside the main reaction.
• Purification of the product may results in loss of some product.
• Percentage yield = how much product was actually made compared with the amount of product that was expected.
• Theoretical yield = the maximum mass of product expected from the reaction, calculated using reacting masses.
• Actual yield = the mass of product that is actually obtained from the real chemical reaction.
• Limiting reagent = the reactant that is not in excess, and is completely used up first.
• Atom economy = a measure of how well atoms have been utilised.
• Reactions with high atom economies:
- Produce a large proportion of desired products and a few unwanted waste
products.
- Are important for sustainability, as they make the best use of natural resources.
Sustainability:
• The process uses reactants that are readily available, carbon from coal and steam from water. Energy will be needed to produce the
steam, but costs for obtaining starting materials are low.
• Other reactions may have a much larger atom economy, but poor % yield. Efficiency depends on both factors.
4.1 Acids, bases, and neutralisation
• All acids have H in their formula – when dissolved in water, an acids releases hydrogen ions as protons, H+, into the solution.
• Acid = proton donor – releases H+ ions in aq. solution.
• Base= proton acceptor – accepts H+ ions in aq. solution.
• Salt = a chemical compound formed from an acid when an H+ ion from acid has been replaced by a metal ion or another positive ion.
• Strong acid = releases all its H atoms into solution as H+ ions and completely dissolves in aqueous solution.
• E.g. HCl (aq) + aq → H+ (aq) + Cl- (aq)
• Weak acid = releases a small proportion of its available H atoms into solution as H+ ions. It partially dissociates in aqueous solution.
• E.g. CH3COOH (aq) + aq ⇌ H+ (aq) + CH3COO- (aq)
• Base = neutralises an acid to form a salt.
• Alkali = a base that dissolves in water, releasing hydroxide ions into the solution, e.g. NaOH (s) + aq → Na+ (aq) + OH- (aq)
• Bases are metal oxides, metal hydroxides, metal carbonates and ammonia.
• In neutralisation of an acids: H+ (aq) ions + base → salt + water.
• The H+ from the acid are replaces by metal or ammonium ions from the base.
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,OCR A-Level Chemistry
Rushali Chawan 13BWD
4.2 Acid-base titrations
• Titration = a technique used to accurately measure the volume of one solution that reacts exactly with another solution.
• They can be used for: (1) finding the concentration of a solution (2) identifying unknown chemicals (3) finding purity of a substance
Preparing a standard solution:
• Standard solution = a solution of known concentration.
• Volumetric flask = used to make up a standard solution very accurately.
• Can’t make a standard solution from:
- Sulfuric acid: can’t get it as a solid (to dissolve in water).
- Sodium hydroxide: absorbs water = less NaOH (as part of it would be water).
- Lead (II) sulfate: absorbs water = less PbSO4 (as part of it would be water).
Steps:
1. Weigh the solid → weigh the boat and then weight boat + solid
2. Dissolve the soil in a beaker, using less distilled water than needed to fill the volumetric flask to the mark.
3. Transfer this solution to a volumetric flask using a funnel. Rinse the funnel and the last traces of the solution into the flask with
distilled water.
4. Fill the flask to the graduation line by adding distilled water drop-by-drop until the bottom of the meniscus lines with the mark.
5. Add a stopper and invert the volumetric flask several times to mix the solution, to get it consistent.
Acid-base titrations:
Apparatus:
• A solution of acid is titrated against a solution of base using a pipette and burette:
• A burette reading is recorded to the nearest half division, with the bottom of the meniscus on a mark/between 2 marks.
• Each burette reading is measured to the nearest ±0.05cm3, so the reading always has 2 D.P, either a 5 or 0.
Method:
1. Add a measured volume of one solution to a conical flask using a pipette.
2. Add the other solution to a burette, and record the initial burette reading to the nearest 0.05cm3.
- When filling the burette, run excess solution out through the top to remove air bubbles.
- If there is an air bubble, air could be released during the titration = error in the titre.
3. Add a few drops of an indicator to the conical flask solution.
4. Put the burette solution into the conical flask solution – swirl to mix the two solutions.
- The indicator changes colour at the end point of the titration, which is used to indicate the volume of the 1st solution
that exactly reacts with the volume of the 2nd solution.
5. Record the final burette reading – read from bottom of meniscus, at eye level. Titre = the volume of solution added from the burette
(final – initial burette reading).
6. A quick, trial titration is carried out first to find the approximate titre.
7. The titration is then repeated accurately, adding the solution dropwise as the end point is approached. Repeat titrations until 2
accurate titres are concordant – agree to within 0.10cm3.
The mean titre:
• To work out the mean titre, use only the closest, accurate titres.
• By repeating titres until 2 agree within 0.10cm3, you can reject inaccurate titres.
• If you included ALL titres = you lose the accuracy of the titration technique.
5.1 Electron structure
• Shells are energy levels – the energy increases as the shell number increases.
• Principal quantum number, n = the shell/energy level number.
• Shells are made of atomic orbitals.
• Atomic orbitals = a region around the nucleus that can hold up to 2 electrons, with opposite
spins. It is a 3D region of space where there is a high probability of finding an electron.
• The greater the shell number n, the greater the radius of its s-orbital.
• The greater the shell number n, the further the p-orbital is from the nucleus.
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,OCR A-Level Chemistry
Rushali Chawan 13BWD
5.1 Electron structure
Filling up orbitals:
1. Orbitals fill in order of increasing energy.
• The sub-shells that make up shells have slightly different energy levels.
• Within each shell, the new type of sub-shell added has a higher energy.
• E.g. in the 2nd shell, the 2p sub-shell is the new type, so has a higher energy than the 2s sub-shell.
• The highest energy level in the 3rd shell overlaps with the lowest energy level in the 4th shell.
• The order of filling is The Aufbau Principle: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f.
2. Electrons pair with opposite spins.
• Each orbital can hold up to 2 electrons.
• We represent different orbitals as boxes and use arrows to show the electrons.
• Electrons are -/ive charged and repel one another.
• Electrons have a property, called spin – either up or down, shown by an arrow: or
• The two electrons in the same orbital must have opposite spins.
3. Orbitals with the same energy are occupied singly first.
• Within a sub-shell, the orbitals have the same energy.
• One electron occupies each orbital before paring starts.
• This prevents any repulsion between paired electrons until there is no further orbital available at the same energy.
Electron configuration of atoms:
• Electron configuration = shows how sub-shells are occupied by electrons. So, Kr = 1s22s22p63s23p63d104s24p6
Highest energy sub-shell:
Blocks: the elements in the periodic table can be divided in to blocks corresponding to their highest energy subshell.
Exceptions:
Cr = 1s2 2s2 2p6 3s2 3p6 4s1 3d5
NOT 1s2 2s2 2p6 3s2 3p6 4s2 3d4
Cu = 1s2 2s2 2p6 3s2 3p6 4s1 3d10
NOT 1s2 2s2 2p6 3s2 3p6 4s2 3d9
Shorthand electron configurations:
Electron configurations can be expressed more simply. This is very useful for emphasizing similarities in outer shell configurations.
1. Identify the previous noble gas in the periodic table e.g. for K, it would be Ar.
2. Put the noble gas symbol in brackets [Ar].
3. Add the outer shell configuration for the atom (or ion) → [Ar] 4s1.
Electron configuration of ions:
• The highest energy electrons are lost when an ion is formed.
• EXCEPTION: 4s electrons are lost before 3d as once 4s and 3d are occupied, 4s moves above 3d.
• Ions of d-block elements: the 4s sub-shell fills AND empties before the 3d sub-shell.
5.2 Ionic bonding and structure
Ionic bonding = the electrostatic attraction between positive and negative ions. It holds together cations (+) and anions (-) in ionic
compounds.
• Outer-shell electrons from a metal atom are transferred to the outer shell of a non-metal atom.
• Positive and negative ions are formed.
• The ions formed often have outer shells with the same electron conjuration as the nearest noble gas.
• Structure of ionic compounds: each ion attracts oppositely charged ions in all directions, forming a giant ionic lattice.
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,OCR A-Level Chemistry
Rushali Chawan 13BWD
5.2 Ionic bonding and structure
Structure of ionic compounds:
1. Electrical conductivity:
• Ionic compounds only conduct electricity when molten OR dissolved in water: solid ionic lattice breaks down + ions are free to
move and carry charge.
• When solid, ions are fixed in position and can’t move.
2. Melting and boiling points:
• Most ionic compounds are solid at RT – they have high melting/boiling points.
• High temperatures are needed to provide a large quantity of energy to overcome the strong electrostatic forces of attraction
between oppositely charged ions in the giant ionic lattice.
• Ionic charge – the greater the charge = the stronger the electrostatic force = the higher the MP/BP.
• Ionic radius – the smaller the ionic radius = the stronger the electrostatic force = the higher the MP/BP.
3. Solubility:
• Many ionic compounds dissolve in polar solvents (such as water).
• Polar water molecules break down the lattice and surround each ion in solution.
• If the ions have large charges = the attraction between the ions may be too strong for the water to break down = not soluble
in water.
• Solubility requires 2 processes: (1) ionic lattice must be broken down + (2) water molecules must attract and surround the
ions.
• So, solubility of an ionic compounds in water depends on their relative strengths of the attractions within the giant ionic lattice
+ the attraction between ions and water molecules.
• As ionic charge increases, solubility decreases.
4. Strength:
• Ionic substances are brittle – if you move a layer of ions, you get ions of the same charge next to one another = layers repel
each other + the crystal breaks up.
5.3 Covalent bonding
Covalent compounds and molecules:
• Covalent bond = the strong electrostatic force between a shared pair of electrons and the nuclei of the bonded atoms.
• There are CBs in: non-metallic elements (H2 and O2) + compounds of non-metallic elements (H2O and CO2) + polyatomic ions (NH4+)
• Atoms are bonded together in a single unit – a small molecule, a giant covalent structure or a charged polyatomic ion.
Orbital overlap:
• Covalent bond = the overlap of atomic orbitals, containing 1 electron, to give a shared pair of electrons.
• The shared pair of electrons is attracted to the nuclei of both the bonding atoms.
• The bonded atoms often have outer shells with the same electron structure as the nearest noble gas.
A covalent bond is localised:
• In a covalent bond, the attraction is localised, acting between the shared pair of electrons and the nuclei of the 2 bonded atoms –
the results is a small unit, a molecule, with 2 or more atoms.
- Molecule = the smallest part of a covalent compound that can exist whilst retaining the chemical properties of the compound.
Dot-and-cross diagrams:
• Displayed formula = shows the relative positioning of atoms and the bonds between them as lines.
• Lone pairs = paired electrons that are not shared.
BF3:
• Boron has the electron configuration: 1s22s22p1 – only 3 outer-shell electrons can be paired.
- Unpaired electrons pair up.
• The maximum number of electrons that can pair up is = to the number of electrons in the outer shell.
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, OCR A-Level Chemistry
Rushali Chawan 13BWD
5.3 Covalent bonding
Phosphorus, sulfur and chlorine:
• For P2 elements, the n = 2 outer shell can hold 8 electrons.
• For phosphorus, sulfur and fluorine, the n = 3 outer-shell can hold 18 electrons.
Example – sulfur hexafluoride, SF6:
• In SF6, 6 unpaired electrons from S are paired.
• The outer shell of S now has 12 electrons – far more than the nearest noble gas, Ar.
• This is called expansion off the octet – only possible from n = 3 shell, as a d-sub-shell becomes available.
Multiple covalent bond = occurs when 2 atoms share more than one pair of electrons.
Double covalent bonds:
• In a double bond – the electrostatic attraction is between 2 shared pair of electrons and the nuclei of the bonding atoms.
• All atoms have 8 electrons in their outer shell, and the electron structure of the nearest noble gas.
Triple covalent bonds:
• In a triple bond – the electrostatic attraction is between 3 shared pairs of electrons and the nuclei of the bonding atoms.
• Again, all atoms have the electron structure of the nearest noble gas.
Dative (co-ordinate) covalent bonds:
• Dative bond = a covalent bonds where the shared pair of electrons comes from one of the bonding atoms only.
- The shared electron pair was originally a lone pair of electrons on 1 of the bonded atoms.
Example – formation of ammonium ion, NH4+:
• Average bond enthalpy = a measurement of covalent bond strength.
• The larger the average bond enthalpy, the stronger the covalent bond.
6.1 Shapes of molecules and ions
Electron-pair repulsion theory:
• Electrons repel each other – as they have negative charge.
• The electron pairs repel one another – so that they are arranged as far as possible.
• The arrangement of electron pairs minimises repulsion – holds the bonded atoms in a definite shape.
• Different numbers of electron pairs = different shapes.
• The greater the number of electron pairs, the smaller the bond angle.
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