A-Level Chemistry

• Isotopes are atoms with the same number of protons and electrons but a different number of neutrons.
• Relative isotopic mass is the mass of an atom of an isotope compared with 1/12 of the mass of an atom of Carbon-12
• Relative atomic mass is the average weighted mean of an atom of an element compared with 1/12 of the mass of Carbon-12
• Avoagadro's number is 6.02 x 1023
• Molar mass is the mass per mole of substance
• Emprical formula is the simplest whole number ratio of atoms of each element present in a compund
• Molecular formula is the actual number of atoms of each element in a molecule
• Empirical formula = (mass/% composition)/molar mass (do this on both sides); then divide by the smallest numer of moles to find the ratio

• Things to construct a chemical eqaution

1. Group 7, oxygen, nitrogen, and hydrogen are diatomic

2. State symobols should always be included in the answer even if they aren't asked for

3. Answers should always be balanced even if they don't ask for it

• Number of moles = mass/molar mass (solid)

• Number of moles = volume/24 (gas)

• Number of moles = conc x volume
• Concentarion - amount of solute, in mol, per 1dm3, of solution

• Acid is a proton donor
• Hydrochloric Acid- HCl
• Sulphuric Acid - H2SO4
• Nitric Acid - HNO3
• Base - proton acceptor and they neuralise acids
• Sodium Hydroxide - NaOH
• Potassium Hydroxide - KOH
• Ammonia - NH3
• Alkali's are special bases that dissolve in water forming aqueous hydroxide ions
• Hydroxide ions from alkali's neutralise the protons from the acid's forming water

• Acid + Carbonate = Salt + H2O + CO2
• Acid + Base = Salt + Water
• Acid + Alkali = Salt + Water
• Acid + Aqueous Ammonia = Ammonium salt
• NH3 + HNO3 = NH4NO3
• Anhydrous - substance containing no water molecules
• Hydrated - crystalline compund containing water molecules
• Water of Crystallisation - water molecules that form an essential part of the crystalline structure of a compound
• How to calculate the formula of a hydrated salt:
• 1. Mass of hydrated - Mass of anhydrous = Mass of H2O
• 2. Calculate number of moles of anhydrous substance
• 3. Calculate number of moles of water
• 4. Look at molar ratio and divide by the smallest number to find the formula of hydrated salt

• Uncombined element = 0
• Combined oxygen = -2
• Combined hydrogen = +1
• Simple ion = same as charge on ion
• Combined fluorine = -1
• Oxidation - loss of electrons
• Reduction - gain of electrons
• Oxidation number increases, oxidation has occured
• Oxidation number decreases, reductions has occured
• Metals generally form ions by losing electrons with an increase in oxidation number to form positive ions
• Non-metals generally react by gaining electrons with a decrease in oxidation number to form negative ions

• First ionisation energy - energy required to remove one electron for each atom in one mole of gaseous atoms
• Successive ionisation energy - energy required to remove each electron in turn
• Nuclear charge - the greater the nuclear charge, the greater the attractive force of the outer electrons; therefore, the more energy needed to remove an electron from the outer shell (this increases across a period)
• Electron shielidng - more inner electron shells shield the nuclear charge fom the outer electron, so the electron is easier to remove (sheilding increases down a group)
• Atomic radius - in larger atoms, the outer electrons are further from the nucleus due to the number of shells; this means that nuclear attraction is lower so the energy needed to remove an electron is much lower (this increases down a group)
• Group 1,2, and 3 elemnts have low ionsiation energies

• 1st shell - 2
• 2nd shell- 8
• 3rd shell - 18
• 4th shell - 32
• S orbitals are spherical and can hold two electrons in total
• P orbitals are dumb-bell shaped and can hold two electrons in each orbital
• In an S subshell, there is only one orbital which can hold two electrons of opposite spins
• In a P subshell, there are 3 orbitals, can hold 6 electrons in total
• In a D subshell there are 5 orbitals, can hold 10 electrons in total
•  In an F subshell there are 7 orbitals, can hold 14 electrons in total
• 1s, 2s, 2p, 3s, 3p, 4s, 3d,4p,4d, 4f
• S block is groups 1, 2, hydrogen, and helium
• P block is groups 3-8 minus helium
• D block is transition metals

• Ionic bond - electrostatic attraction between oppositely charged ions
• BE, B,C, and Sl don't normally form ions and too much energy is needed to transfer the outer-shell electrons to form ions
• Nitrate - NO3-
• Carbonate - CO3 2-
• Sulphate - SO4 2-
• Ammonium - NH4+ 
• Covalent bond - a shared pair of electrons
• Dative covalent bond - two electrons derive from the same atom

• The shape of a simple molecule if determined by repulsion between electron pairs surrounding the central atom
• Lone pairs of electrons repel more than bonded pairs
• BF3 (trigonal planar) - 120 degrees
• CH4 & NH4+ (tetrahedral) - 109.5 degrees
• SF6 (octahedral) - 90 degrees
• NH3 (pyramidal) - 107 degrees
• H2O (non-linear) - 104.5 degrees
• CO2 (linear) - 180 degrees
• Electronegativity is the ability of an atom to attract the bonding electrons of a covalent bond
• If an atom is more electronegative than the other atom, it will have a greater attraction for the bonding pair of electons than the other atom, so the bonding electrons will be closer to it. This creates a small charge difference which is a permanent dipole-dipole force

• Intermolecular forces - based on permanent dipoles, as in hydrogen chloride, and induced dipoles (london dispersion forces), as in the noble gases
• Polar molecules have permanent dipoles. The permanent diopole of one molecule attracts another permanent dipole to form a week permanent diople-dipole force
• 1.The movement of electrons in shells unbalances the distribution of the charge in the electron shells
• 2. An instantaneous dipole is formed due to the unbalanced distribution of electrons
• 3. The instantaneous dipole induces a dipole in neighbouring molecules, which then turn induce firther molecules
• 4. The induced dipoles attract each other, forming weak intermolecular bonds
• A hydrogen bond is a strong dipole-dipole intercation (hydorgen bonding)
• Metallic bonding - attraction of positive ions to delocalised electrons

• Giant ionic lattices have strong ionic bondong e.g. NaCl
• Formed by the attraction of oppositley charged ions
• 1. Each ion is surrounded by oppositely charged ions
• 2. The ions attract each other to form a giant ionic lattice
• Giant covalent lattices e.g. diamond, and graphite
• A 3D structure of atoms bonded together by strong intramolecular covalent bonds
• Giant metallic lattices:
• 1. Contains ionised atoms
• 2. Positive ions occupy fixed positions in a lattice
• 3. Delocalised outer shell electrons
• 4. The delocalised electrons are spread throughout the metallic structure
• 5. The electrons can move within the structure
• 6. The charges balance over the whole structre
• Simple molecular lattices:
• A 3D structure of molecules bonded together by weak intermolecular forces
• 1. The molecules are held together by weak forces between them
• 2. The atoms within each molecules are bonded strong,y together by covalent bonds
• When a simple molecular structure is broekn, it is the weak London Dispersion Forces and not the covalent bonds broken

• Metallic structures:
• 1. High melting & boiling point (lots of energy needed to break the bonds)
• 2. Good electrical conductivity (deloclaised electrons carry charge)
• Malleable & ductile (delocalised electrons so layers can slide over each other)
• Giant ionic compounds:
• 1. High melting and boiling point (lots of energy required to break the bonds)
• 2. Doesn't conduct when solid (ions are fixed and can't move)
• 3. Conducts when molten (ions are free to move)
• 4. Can dissolve in polar solvents (H2O molecules break down the lattice by surrounding each ion)
• Simple molecular structures:
• 1. Low melting and boiling points (weak LDF's so less energy needed to break the bonds)
• 2. Doesn't conduct (no charged particles)
• 3. Soluble in non-polar substances (LDF's form between the SMS and the non-polar substance which weakes the lattice structure)
• Giant covalent  structures:
• 1. High melting and boiling point (lots of energ required to break the bonds)
• 2. Doesn't conduct (no free charged particles)
• 3. Insoulable in both polar and non-polar substances (covalent bonds are too great to be broken by both solvents)

• Graphite:

• 1. Good conductivity (delocalised electrons between layers are free to move parallel to the layers where a voltage is applied)

• 2. Soft (bonding in each layer is strong but weak LDF's between layers allow them to slide easily)

• The Periodic Table is arranged by:
• 1. Increasing atomic number
• 2. Periods showing repeating trends in physical and chemical properties
• 3. Groups having similar physical and chemical properties, due to having the same number of electrons in the outer shell
• Atoms of the same elements in a group have similar outer shell electron configurations, resulting in similar properties
•  Periodicity - trends or recurring variations in element properties with increasing atomic number

• The variation of the first ionistaion energies of elements are caused by:
• 1. A general increase across a period, in terms of nuclear charge
• a) as you go across a period, the number of protons increases, so there is more energy is needed to remove electrons from the outer shell
• b) across a period, electrons are added to the outer shell so it is drawn inwards slightly, but the electrons shielding hardly changes
• 2. A decrease down a group in terms of increasing atomic radius and increasing electron shielding outweighing increasing nuclear charge
• a) down a group, the number of shells increases. This increases the atomic radius so therefore the attraction between the nucleus and the electrons is less. So ionistaion enegy decreases
• b) down a group, there are also more inner shells so the shielding effect of the outer electrons increases, so ther is, again, less attraction between the nucleus and the energy
• c) the nuclear charge also increases, but this doesn't have much of an effect on the ionisation energy

• The variation in electron configurations, atomic radii, m&b points are caused by:
• 1. Period 2 elements all ahve two outer shells, whereas period 3 elements all have three outer shells
• 2. Period three has an extra electron shell than period 2, so the atomic radius is larger
• 3. Boiling points in group 1-4 in both periods are fairly high, this is because group 1-4 have strong forces holding the structures together
• 4. Between group 4 and 6, there is a sharp decrease in BP, this marks the change between the strong forces between atoms to the weak forces between molecules
• 5. In group 5-8, the BP's are all comparitively low due to the weak forces
• Variations in M&B P's in terms of structure and bonding happen because:
• 1. In period 2, the first two groups are metals so the bonding is metallic, which is a strong force between the molecules
• 2. Group 3&4 are then the giant covalent structures, so still have strong forces between the atoms, so the BP remains high
• 3. After group 4 in period 2, the forces chnage to weak LDF's as the elements are no longer metals
• 4. This causes a sharp decrease in M&BP as the elements as less energy is needed to break the bonds between the molecules
• 5. This is the same in period 3, apart from group 3 is also metallic bonding

• Group 2 elements react vigorusly with oxygen. A redox reaction occurs where an ionic oxide is produced with the formula MO (where M = group 2 metal)
• Group 2 elements react with water to form hydroxides, with the general formula: M(OH)2 and hydrogen gas is also formed
• Group 2 elements are both reactive metals and strong reducing agents
• 1. Group 2 elements are oxidised in reactions. Each atom loses two electrons from its outer S subshell for form a 2+ ion
• 2. Reactivity increases down the group, due to increased atomic radius and nuclear shieling the forces of attraction between the electrons and the nucleus is lower, so it is easier to lose electrons down the group

• Group 2 oxides react with water to form a solution of metal hydroxide
• The typical pH of these solutions is between 10-12
• Group 2 carbonates are decomposed by heat (thermal decomposition), forming a metal oxide and carbon dioxide
• Carbonates become more difficult to decompose with heat as you go down a group
• Group 2 elements are more reactive as you go down the group
• Group 2 carbonates decompose at higher temperatures down the group
• Hydroxides are more soluable in water, and the resulting solutions become more alkaline down the group
• Calcium hydroxide (Ca(OH)2) is used as "lime" by farmers and gardeners to neutralise acidic soils, using too much makes the soil too alkaline
• Magnesium hydroxide (Mg(OH)2) is used as "milk of magnesia" to relieve indigestion. It neutralises excess acid in the stomach

• As you move down group 7, the boiling point increases because:
• 1. The number of electrons increases, leading to LDF's between molecules
• 2. Therefore, more energy is needed to break the bonds, so hence a higher boiling point
• Cl2 in water = pale green / Cl2 in cyclohexane = pale green
• Br2 in water = orange / Br2 in cyclohexane = orange
• I2 in water = brown / I2 in cyclohexane = violet
• Cl2 + 2Br-  = 2Cl- + Br2
• Cl2 + 2I-  = 2Cl- + I2
• Br2 + 2I-  = 2Br- + I2
• As you go down group 7, they become less reactive because:
• 1. Atomic radius increases
• 2. Shielding increases
• 3. The nuclear attraction between the electrons and the nucleus increases, making it harder for the outer shell to gain an electron into the p-subshell
• Cl2 +H2O = HClO + HCl
• Cl2 + 2NaOH = NaCl + NaClO + H2O
• Disproportionation - both oxidised and reduced

• Chlorine makes water safer to drink by sterilising it through killing bacteria
• However, chlorine reacts with organic matter to form chloronated hydrocarbons, which may cause cancer
• Chlorine is also toxic

• An unknown halide is dissolved in water
• Aqeuous silver nitrate (AgNO3) is added
• The Ag+ ions react with the halide ions to form coloured precipitates
• If unsure abaout the colour, adding aqeous ammonia (NH3) may help distinguish as different precipitates have different solubility in the ammonia
• Ag+ + Cl-   = AgCl
• Ag+ + Br-  = AgBr
• Ag+ + I-  = AgI
• These are called precipiatation reactions (they take place in aqueous solutions whne aqueous ions react together to form solid precipitate)