Topic 4 Inorganic chemistry and the periodic table

The atomic radius increases as we go down Group 2. This is as extra shells are added as you go down All group 2 metals lose 2 electrons to form +2 ions when they react. All group 2 metals have electron configurations that end in s2

1st IE decreases as we go down group 2. There is more shielding hence weaker attraction between the nucleus and outer electrons. The outer electrons are also further from the nucleus which weakens the attraction. We do have an increase in the number of protons as we go down the group. However, the shielding effect overrides an increase in + charge

React with water to form metal hydroxides Sr(s) + 2H2O(l) —–> Sr(OH)2(aq) + H2(g) Reactivity increases with water as you go down the group. (There is no reaction with Be) The reason why is that the atom gets larger and the electron is further from the nucleus. Easier to remove and hence more reactive. There is more shielding Magnesium reacts slowly with cold water but more rigorously with steam. This produces Magnesium Oxide(MgO) instead of hydroxide

They react with oxygen to form metal oxides 2Mg(s) + O2(g) —–> 2MgO(s) The oxidation number of Mg increases from 0 to +2. Magnesium is being oxidised The oxidation number of O decreases from 0 to -2. Oxygen is being reduced Group 2 oxides are white solids

They react with chlorine to form metal chlorides. Mg(s) + Cl2(g) —-> MgCl2(s) The oxidation number of Mg increases from 0 to +2. Magnesium is being oxidised The oxidation number of chlorine has been reduced from 0 to -1. Chlorine is being reduced.

They react with water to form bases. Alkaline solutions formed. SrO(s) + H2O(l) —-> Sr(OH)2(aq) SrO(s) + H2O9l) —-> Sr2+(aq) + 2OH-(aq) Oxides react readily with water to make hydroxides which dissociate to form OH- ions. (Magnesium oxide reacts very slowly and the hydroxide barely dissolves. Beryllium oxide doesn’t react with water at all and the hydroxide is insoluble) They become more strongly alkaline as we go down the group as the hydroxides become more soluble

They can neutralise acids How oxides react with acids : Cao(s) + 2HCl(aq) —-> CaCl2(aq) + H2O(l) How hydroxides react with acids : Ca(OH)2(s) + 2HCl(aq) —> CaCl2(aq) + 2H2O(l)

As a general rule, if the -ion has a double charge, they become less soluble as we go down a group Generally, if the - ion has a single - charge they become more soluble down the group

Group 2 carbonates and nitrates can decompose upon heating Carbonates break down into metal oxides and carbon dioxide. This is via thermal decomposition CaCO3(s) —-> CaO(s) + Co2(g) Nitrates break down into metal oxides, nitrogen dioxide and oxygen via thermal decomposition 2Ca(No3)2(s) —> 2CaO(s) +4NO2(g) + O2 Carbonates/Nitrates become more thermally stable as we go down group 2. The carbonate/nitrate ion has a large electron cloud that can be distorted when nearby + group 2 metal ions All group 2 metal ions have a +2 charge however they become larger as we go down the group meaning the charge is spread out over a larger area. They have a lower charge density Mg2+ has a high charge density and distorts the electron cloud in carbonates/nitrates ions more than Ba2+ which has a lower charge density. The less distortion the more stable the carbonate is.

Group 1 carbonates can also decompose upon heating and are more thermally stable than group 2 compounds. Carbonates are thermally stable under a Bunsen flame. Except for lithium carbonate which forms an oxide and CO2. Lo2Co3(s) —-> Li2O(s) + CO2(g) Nitrates break down nitrites and oxygen (except LiNO3 decomposes to form Li2O + NO2+ O2) 2KNO3(s) —-> 2KNO2(s) + O2(g) No3- is a Nitrate ion NO2- is a nitrite ion

Nitrates: Measure how long it takes a specific amount of oxygen to be produced. Using a gas syringe or the amount needed to relight a glowing splint The length of time it takes to produce a specific amount of NO2. It is a brown gas so it can be easily observed. It is toxic so it must be done in a fume cupboard Carbonates: The length of time it takes until a specific amount of CO2 is produced. Co2 turns limewater cloudy so the quicker this turns cloudy the more carbon dioxide is produced. Could use a gas syringe too.

Test for + ions (cations) in a compound using flame tests Electrons in the shells move to higher energy levels as they absorb energy from the flame. When they drop back down to lower energy levels light is released. Different colours are produced as the difference in energy levels determines the wavelength of light released.

Dip the nichrome wire in concentrated HCl. Dip into the sample. Place the loop into the blue flame and observe the colour Calcium - Dark red Rubidium - Red Lithium - Crimson Strontium - Crimson Sodium - Yellow-orange Barium - Green Caesium - Blue Potassium - Lilac

They make up group 7 Fluorine is a pale yellow gas Chlorine is a pale green gas Bromine is a brown-orange liquid Iodine is a grey solid The Boiling points increase as we go down the group. This is because the London forces increase due to increasing size and relative mass of the atoms. The physical state goes from gas at the top of group 7 to solid at the bottom Electronegativity decreases as we go down the group. Electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond. The atom gets larger and the distance between the + nucleus and bonding electrons increases. There is also more shielding

Adding hexane can make it easier to observe colour changes. When undertaking halogen displacement reactions we can add an organic solvent like hexane to see the colour change easily TOBA The halogen present will dissolve readily in the organic solvent which forms a layer above the aqueous layer. A coloured band will appear. Organic is on top of aqueous More reactive halogen will displace less reactive halide ions Reactivity in halogens decreases as we go down group 7. For a reaction to occur an electron is gained. Atoms with a smaller radius attract electrons better than larger atoms. Halogens are less oxidising as we go down the group. We can show this by reacting halogens with halide ions

KCL reacts with Cl2 - No reaction Aq is colourless Org is colourless KBr reacts with Cl2 - Aq is yellow Org is orange KI reacts with Cl2 - Aq is brown Org is purple KCl reacts with Br2 - No reaction Aq is yellow Org is orange KBr reacts with Br2 - No reaction Aq is yellow and Org is orange KI reacts with Br2 - Aq is brown Org is purple KCl reacts with I2 - No reaction Brown and purple KBr reacts with I2 - No reaction brown and purple KI reacts with I2 - No reaction Brown and Purple

They react to form metal halides Mg + Cl2 —> MgCl2 Oxidation number of Mg increases from 0 to + 2. Mg is being oxidised. Mg —> Mg2+ + 2e- Oxidation number of cl2 is reduced from 0 to -1. Cl2 + 2e- —> 2cl- 2Li(s) + Cl2(g) —-> 2LiCl(s) Oxidation number of Li increases from 0 to +1. Lithium is being oxidised Oxidation number of cl2 decreases from 0 to -1. Chlorine is being reduced

Halogens react with COLD alkalis in disproportionation ions Br2 + 2NaOH —> NaOBr + NaBr + H20 Br2 + 2OH- —> OBr- + Br- + H2O Oxidation number of Br is 0 to +1 to -1 Halogens react with HOT alkalis in disproportionation reactions 3Br2 + 6OH- —> BrO3- + 5Br- + 3H2O Oxidation number of X goes from 0 to +5 to -1

Mixing chlorine and sodium hydroxide will form sodium chlorate solution - AKA bleach 2NaOH(aq) + Cl2(g) —-> NaClO(aq) + NaCl(aq) + H20(l) Chlorine goes from 0 to +1 to -1 Uses of sodium chlorate: Treating water Bleaching paper and fabrics Cleaning agents

Adding chlorine to water can kill bacteria. It will produce chlorate ions (Clo-) which kills bacteria. Useful in drinking water and pools H2O + CL2 <======> HCL + HCLO Chlorine has been simultaneously reduced and oxidised. We call this a disproportionation reaction. HCLO + H2O <======> CLO- + H3O+

Halide ions lose an electron in reactions as they are reducing agents. As we go down the group the ionic radius increases The distance between the nucleus and outer electrons becomes larger and there is more shielding. The attractive force gets weaker. The outer electron is lost more readily and this is the reason why I- is a more powerful reducing agent than F- There are 2 tests to prove this trend: 1 - Reaction with sulfuric acid 2- Reaction with silver nitrate solution

Oxidation number of sulfur in NaHSO4 SO2 S H2S + 6 + 4 0 -2 Point A With NaCl : H2SO4 + NaCl —–> NaHSO4 + HCl With NaBr : H2SO4 + NaCl —–> NaHSO4 + HBr With NaI : H2SO4 + NaCl —–> NaHSO4 + HCI Point B in addition to A With NaBr ( same reaction with NaI) : 2Br- —-> Br2 + 2e- (Br- ions oxidised) H2SO4 + 2H+ +2e- —–> SO2 + 2H2O (S being reduced) Overall ionic is H2SO4 + 3H+ + 2Br- —-> Br2 + SO2+ 2H2O (Orange vapour of Br2 produced) Point C in addition to point A and B 6I- —-> 3I2 + 6e- (I- ions oxidised) H2SO4 + 6H+ + 6e- —-> S + 4H2O 9S being reduced) Overall ionic - H2SO4 + 6H+ + 6I- —-> 3I2 + S + 4H2O (Yellow solid of S produced)

They are acidic They are gases that dissolve in water to form acidic solutions. They react with water in the air to form white misty fumes. Hydrogen halides react with ammonia gas to make white fumes of ammonium halides NH3 + HCl —> NH4Cl

Tests for halides using silver nitrate and then confirm with ammonia solution. Add dilute HNO3 then silver nitrate solution AGNO3. The colour of the precipitate will help you to identify the halide ion. The only reason we add nitric acid is for it to react with any anions other than halides otherwise you could get a false result Chloride ions - White precipitate Bromide ions - Cream precipitate Iodide ions - Yellow precipitate Further test - Add ammonia NH3 Cl- white precipitate dissolves in dilute NH3 Br- cream precipitate dissolves in concentrated NH3 I- yellow precipitate is insoluble in concentrated NH3