Edexcel Chemistry - Topic 4: Inorganic Chemistry I - Group 2 and 7

• Atomic radii increases down the group as there are more shells of electrons
• Ionic radii increases down the group as theare more shells increase shielding
• First ionisation energy decreases down the group as their is more shielding and a larger atomic radius so the nuclear attraction between the nucleus and valence electrons are weaker
• Melting points decreases down the group as the ionic size increases weakening the metallic bond
• Reactivity increases down the group as the valence electrons are more easily lost forming positive ions because of the increased shielding
• Electronegativity decreases down the group as the nuclear attraction becomes weaker

Reactions with Oxygen

• Magnesium burns in oxygen with a bright white flame
• 2Mg + O2 --> 2MgO
• MgO is a white solid with a high melting point because of ionic bonding
• Mg slowly reacts with oxygen without a flame so can develop a layer of MgO on its surface
• MgO can be detected by reacting it with hydrochloric acid. If it is present, it will react with the acid to form water and not hydrogen (which is detected using the squeaky pop test)
• Mg + 2HCl --> MgCl2 + H2
• MgO + 2HCl --> MgCl2 + H2O

Reactions with Chlorine

• Group 2 metals react with chlorine to form a salt
• Mg + Cl2 --> MgCl2

Steam

• Mg burns in steam with a bright white flame to produce MgO and hydrogen
• Mg(s) + H2O(g) --> MgO(s) + H2(g)

Warm Water

• Mg reacts with warm water to produce Magnesium Hydroxide
• The reaction is slower and there is no flame
• Mg + 2H2O --> Mg(OH)2 + H2

Cold Water

• The other Group 2 metals react with cold water to form hydroxides
• Sr + 2H2O --> Sr(OH)2 + H2
• The hydroxides formed make the water alkaline
• Observations include fizzing, metal dissolving and the solution heating up (all of which increase as you go down the group). Calcium also forms a white precipitate

Group 2 Oxides and Water

• React to form hydroxides
• The ionic oxides are basic as the oxide ions accept protons to become hydroxide ions in the reaction (acting as a bronsted lowry base)
• Mg(OH)2 is only slightly soluble in water so fewer free OH- ions form so the pH is lower
• MgO(s) + H2O(l) --> Mg(OH)2(s) - pH 9
• CaO(s) + H2O(l) --> Ca(OH)2(s) - pH 12

Group 2 Oxides and Acids

• React to form salts and water
• CaO(s) + 2HCl(aq) --> CaCl2(aq) + H2O(l)

Group 2 Hydroxides and Acids

• React to form salts and water
• Mg(OH)2(aq) + 2HCl(aq) --> MgCl2(aq) + 2H2O(l)

Hydroxides

• Become more soluble down the group
• When insoluble form white precipitates e.g. Mg(OH)2
• Mg(OH)2 is insoluble and forms a slightly alkaline solution. It is used in medicine (in milk of magnesia) to neutralise excess acid in the stomach and treat constipation. It is safe to use as it is weakly alkaline and preferable to CaCO3 as it doesn't produce CO2
• Ca(OH)2 is slightly soluble and is used to neutralise acids in soils and as a test for CO2 (lime water)

Sulphates

• Become less soluble down the group
• BaSO4 is insoluble in water
• If barium reacts with sulphuric acid it will only react slowly as the insoluble BaSO4 will cover the metal surface acting as a barrier to further attack
• Ba + H2SO4 --> BaSO4 + H2

• Thermal decomposition is the use of heat to break down a reactant into more than one product
• The ease of thermal decomposition decreases down the group
• Group 2 Carbonates are more thermally stable down the group as the larger cations have less of a polarising effect and distort the carbonate ion less. The C-O bond is therefore weakened less so it is harder to break it down.
• Group 1 Carbonates don't decompose except for Lithium as it has a small enough size to have a polarising effect on the Carbonate ion
• CaCO3(s) --> CaO(s) + CO2(g)
• Li2CO3(s) --> Li2O(s) + CO2(g)
• To test ease of decomposition, a known mass of carbonate can be heated in a boiling tube and the gas produced bubbled through limewater. Time for the first permanent cloudiness to appear in the limewater.
  Repeat the experiment with different carbonates using the same number of moles of carbonate/same volume of limewater/same bunsen flame and height of tube above flame

• Group 2 Nitrates decompose to form Group 2 Oxides, Oxygen and Nitrogen Dioxide
• 2Mg(NO3)2 --> 2MgO + 4NO2 + O2
• 2Ca(NO3)2 --> 2CaO + 4NO2 + O2
• NO2 is a brown gas and the white nitrate solid melts to a colourless solution and then resolidifies
• The ease of thermal decomposition and thermal stability of the compounds is the same as the carbonates
• Group 1 Nitrates (except Lithium Nitrate) decompose to give a Nitrate(III) salt and oxygen. Lithium Nitrate decomposes in the same way as Group 2 Nitrates.
• 4LiNO3 --> 2Li2O + 4NO2 + O2
• 2NaNO3 --> 2NaNO2 + O2

• Clean a nichrome (unreactive metal) wire by dipping in conc hydrochloric acid and then heating in the bunsen flame
• Dip the wire in powdered solid, place in the bunsen flame and observe flame colour
• In a flame test, the heat causes the electron to move to a higher energy level. The electron is unstable at the higher energy level and so drops back down. As it drops back down the energy is emitted in the form of visible light energy causing a coloured flame.
• Lithium - Scarlet Red
• Sodium - Yellow
• Potassium - Lilac
• Rubidium - Red
• Caesium - Blue
• Magnesium - Bright White Light
• Calcium - Brick Red
• Strontium - Red
• Barium - Apple Green

Fluorine: pale yellow gas (highly reactive)
Chlorine: greenish gas (poisonous in high concentrations)
Bromine: red liquid (gives off dense brown/orange poisonous fumes)
Iodine: grey solid that sublimes to a purple gas

• Atomic radius increases down the group as there are more shells of electrons
• Melting points increase down the group as there are more electrons so there are stronger Van der Waals Forces between the molecules so more energy is required to break the intermolecular forces
• Electronegativity decreases down the group as the atomic radii increases and there is greater shielding so the nucleus is less able to attract bonding pairs of electrons
• Reactivity decreases down the group as the larger atomic radii and increased shielding mean that the nuclear attraction is weaker so it is harder to attract an electron.

• A halogen that is a stronger oxidising agent will displace a halogen that has a lower oxidising power from one of its compounds
• Oxidising strength decreases down the group. Oxidising agents are electron acceptors
• Chlorine displaces both bromide and iodide ions
• Bromine will displace iodide ions
• The colour of the solution in the test tube shows which free halogen is present in solution
• Chlorine = very pale green/colourless solution
• Bromine = yellow solution
• Iodine = brown solution/black solid
• If an organic solvent is added, the colour of the organic solvent layer in the test tube shows which free halogen is present in solution
• Chlorine = colourless
• Bromine = yellow
• Iodine = purple
• Cl2(aq) + 2Br-(aq) --> 2Cl-(aq) + Br2(aq) = yellow solution/yellow
• Cl2(aq) + 2I-(aq) --> 2Cl-(aq) + I2(aq) = brown solution/purple
• Br2(aq) + 2I-(aq) --> 2Br-(aq) + I2(aq) = brown solution/purple

• Metals can be oxidised by the addition of halogens to form a metal-halide salt
  Br2(l) + 2Na(s) --> 2NaBr(s)
  2Na --> 2Na+ + 2e-
  Br2 + 2e- --> 2Br-
• Chlorine and Bromine can oxidise Fe2+ to Fe3+. Iodine is not strong enough oxidising agent to cause this reaction to occur. Instead Iodine reduces Fe3+ to Fe2+
  Cl2(g) + 2Fe2+(aq) --> 2Cl-(aq) + 2Fe3+(aq)
  2I-(aq) + 2Fe3+(aq) --> I2(aq) + 2Fe2+(aq)
• Chlorine undergoes disproportionation with water, simultaneously reducing and oxidising from 0 to -1 and +1
  Cl2(aq) + H2O(l) --> HClO(aq) + HCl(aq)
• If universal indicator is added, it will first turn red due to the acidity of both products but then will turn colourless as HClO bleaches the colour.
• Chlorine is used in water treatment to kill bacteria. Chlorine is used to treat drinking water and the water in swimming pools.

Cold dilute Sodium Hydroxide

• Cl, Br and I in aqueous solutions react with cold NaOH. The reaction that occurs is a disproportionation reaction. The colour of the halogen solution will fade to colourless

Cl2(aq) + 2NaOH(aq) --> NaCl(aq) + NaClO(aq) + H2O(l)

• The mixture of NaCl and NaClO is used as bleach

Hot dilute Sodium Hydroxide

• With hot alkali, disproportionation also occurs but the halogen that is oxidised goes to a higher oxidation state.

3Cl2(aq) + 6NaOH(aq) --> 5NaCl(aq) + NaClO3(aq) + 3H2O(l)

• In IUPAC nomenclature the chlorine compounds are given their relavent oxidation number in roman numerals e.g. NaClO = Sodium Chlorate(I); NaClO3 = Sodium Chlorate(V)

• The reducing power of halides increases down the group.
• Reducing agents donate electrons. The larger ions have a weaker nuclear attractive force so the valence electrons are more easily lost.
• When NaF or NaCl reacts with conc. sulphuric acid, white steamy fumes of HF or HCl are evolved
  NaCl(s) + H2SO4(l) --> NaHSO4(s) + HCl(g)
• These are acid-base reactions and not redox reactions
• When NaBr reacts with conc. sulphuric acid, the observations include: white steamy fumes of HBr forming, red fumes of bromine and an acidic gas of sulphur dioxide forming
  Acid-Base Step: NaBr(s) + H2SO4(l) --> NaHSO4(s) + HBr(g)
  Redox Step: 2HBr(g) + H2SO4(l) --> Br2(g) + SO2(g) + 2H2O(l)
  Ox equation: 2Br- --> Br2 + 2e-
  Red equation: H2SO4 + 2H+ + 2e- --> SO2 + 2H2O
• Sulphuric acid acts as an oxidising agent in the second redox step of the reaction.

• Iodide ions are the strongest halide reducing agents - they can reduce the sulphur from +6 in H2SO4 to +4 in SO2, 0 in S and -2 in H2S
• Observations include: White steamy fumes of HI, Black solid and purple fumes of Iodine, a colourless acidic gas of sulphuric dioxide, a yellow solid of Sulphur and a colourless gas with a bad egg smell of Hydrogen Sulphide.
  NaI(s) + H2SO4(l) --> NaHSO4(s) + HI(g)
  2HI + H2SO4 --> I2(s) + SO2(g) + 2H2O(l)
  6HI + H2SO4 --> 3I2(s) + S(s) + 4H2O(l)
  8HI + H2SO4 --> 4I2(s) + H2S(g) + 4H2O(l)
• Sulphuric acid acts as the oxidising agent in the three redox stages
  Ox equation: 2I- --> I2 +2e-
  Red equation: H2SO4 + 2H+ + 2e- --> SO2 + 2H2O
  Red equation: H2SO4 + 6H+ + 6e- --> S + 4H2O
  Red equation: H2SO4 + 8H+ + 8e- --> H2S + 4H2O

• The reaction is used to test for halide ions
• The Nitric acid is used to remove any carbonates present that will produce a precipitate that will mask any desired precipitates
• Fluorides produce no precipitate
• Chlorides produce a white precipitate
  Ag+(aq) + Cl-(aq) --> AgCl(s)
• Bromides produce a cream precipitate
  Ag+(aq) + Br-(aq) --> AgBr(s)
• Iodides produce a pale yellow precipitate
  Ag+(aq) + Br-(aq) --> AgI(s)
• Silver halide precipitates can be differentiated by using ammonia solution if the colours look similar
• Silver Chloride dissolves in dilute ammonia
  AgCl(s) + 2NH3(aq) --> [Ag(NH3)2]+(aq) + Cl-(aq)
• Silver Bromide dissolves in concentrated ammonia
  AgBr(s) + 2NH3(aq) --> [Ag(NH3)2]+(aq) + Br-(aq)
• Silver Iodide doesn't react with ammonia.

• Hydrogen halides are made by the reaction of solid sodium halide salts with phosphoric acid
  NaCl(s) + H3PO4(l) --> NaH2PO4(s) + HCl(g)
• Steamy white fumes of the Hydrogen Halides are evolved
• The Steamy fumes of HCl are produced when the HCl meets the air because it dissolves in the moisture in the air
• Hydrogen halides are more dense than air so the delivery tube points downwards to collect the gas
• Phosphoric acid is not an oxidising agent and so doesn't oxidise HBr and HI so there are no extra redox reactions that occur and no other products form.
• Hydrogen halides are all soluble in water and dissolve to form acidic solutions
  HCl(g) + H2O(l) --> H3O+(aq) + Cl-(aq)
• All the hydrogen halides react readily with ammonia to give the white smoke of the ammonium halide. This can be used as a test for Hydrogen Halides
  HCl(g) + NH3(g) --> NH4Cl(s)
  HBr(g) + NH3(g) --> NH4Br(s)

Carbonates and Hydrogencarbonates

• Add any dilute acid and observe effevescence. Bubble the gas through limewater to test for carbon dioxide.
  2HCl + Na2CO3 --> 2NaCl + H2O + CO2
  HCl + NaHCO3 --> NaCl + H2O + CO2

Sulphates

• Addition of Barium Chloride and observe for a white precipitate of BaSO4 will form.
  Ba2+(aq) + SO4 2-(aq) --> BaSO4(s)
• The acid is used to react with any Carbonate impurities that will form a white BaCO3 precipitate and give a false result. Sulphuric acid can't be used as it contains Sulphate ions

Halide Ions

• Addition of silver nitrate to form a precipitate and ammonia to distinguish between precipitates. [see previous]

Ammonium

• Test for ammonium ions by reaction with warm aqueous sodium hydroxide to form ammonia
  NH4+(aq) + OH-(aq) --> NH3(g) + H2O(l)

Metal Ions

• Flame tests can determine whether a metal ion is present

Gases

• Hydrogen - lit splint causes a squeaky pop noise
• Oxygen - relights a glowing splint
• Carbon Dioxide - turns limewater cloudy
• Ammonia - turns red litmus paper blue
• Nitrogen Dioxide - orange-brown gas
• Hydrogen Sulphide - bad egg smell
• Water - turns dry blue cobalt chloride paper pink