Unit 2: Group 2 and 7

Trends in ionisation energy down group

• Ionisation energy decreases because:

1. Atomic radius increases - extra electron shell
2. Shielding from positive nucleus increases
3. Outer electrons are further away - weaker attraction to positive nucleus

Reactivity increases down the group as I.E decreases

How to do a flame test:

• Clean a platinum or nichrome wire by dipping it into HCl and heating it over hot flame 
• Repeat until wire does not produce a colour
• When clean, moisten with some acid and dip into small amount of solid to be tested

Flame colours:

• Lithium - red
• Sodium - yellow
• Potassium - lilac
• Magnesium - No colour
• Calcium - brick red
• Strontium - crimson red
• Barium - green

Origins of colour:

• When heated, electrons gain energy and are promoted to higher energy levels, (unstable)
• Electrons fall back down to original energy levels and give off energy in the form of wavelengths of visible light - this is the colour seen
• Energy released corresponds to the energy gap and therefore a specific colour

Solubility of hydroxides

Increases down the group (Barium hydroxide is soluble so it is poisonous to the body because it would dissolve in the gut)

Solubility of sulfates

Decreases down the group (Barium sulfate is insoluble so it can be used for X-rays of the gut because it is not poisonous)

Factors involved in dissolving:

• 1. Lattice energy - larger attractions means more energy given off and more negative LE
• 2. Enthalpy of hydration - the greater the charge density the easier it is to hydrate with water due to greater attraction of polar molecules to water molecules

1. Oxygen: form metal oxides

2X(s) + O2(g) -> 2MgO(s)

2. Chlorine: form metal chlorides

X(s) + Cl2(g) -> XCl2(s)

3. Water: form metal hydroxides

X(s) + H2O(l) -> X(OH)2(aq) + H2(g)

• Magnesium only reacts rapidly with steam 

Mg(s) + H2O(g) -> MgO(s) + H2(g) 

• Beryllium does not react

Oxides with water: form metal hydroxides

e.g. CaO(s) + H2O(l) -> Ca(OH)2(aq)

• Calcium hydroxide is slaked lime and is used in: water treatment, mortar, plaster etc.
• Observations: fizzing - water vapour is released

Oxides and hydroxides with acids: form salt and water

e.g. oxides: MgO(s) + HCl(aq) -> MgCl2(aq) + H2O(l)

e.g. hydroxides: Ca(OH)2(aq) + 2HNO3(aq) -> Ca(NO3) + 2H2O

• Act as bases and neutralise acids (e.g. HCl, HNO3)

Thermal decomposition: splitting a compound by heating

• The more thermally stable a substance, the more heat required to break it down
• The greater the polarising power of the cation (charge density), the more distortion of the electron cloud and more unstable the anion is. Less heat energy required to break the bond 
• The cationic radius increases down the group so less distortion of the anion is caused and the stability of the anion is greater. 
• Thermal stability increases down the group for carbonates and nitrates 

Carbonates

e.g. CaCo3(s) -> CaO(s) + CO2(g)

• Test: How long it takes for CO2 to be produced - turns limewater cloudy

Nitrates

e.g. Ca(NO3)2(s) -> 2CaO(s) + NO2(g) + O2(g)

• Test: how long it takes for brown gas to be produced - toxic so done in a fume cupboard!

Halogen properties:

• Non-metallic - low melting and boiling points
• Very reactive
• Diatomic covalent molecules - i.e. Cl2, Br2
• Oxidising agents - evident in displacement reactions
• Halides: reducing power increases down group (ability to lose an electron) 
  • Increased shielding as no. of electron shells increases
  • Ions get bigger so electrons are further from the positive nucleus
  • Attraction gets weaker between outer electron and nucleus
• Non-polar - more soluble in organic solvents than water

                                             Fluorine           Chlorine               Bromine                    Iodine

• Physical state & colour:  Yellow gas      Green gas            Red-brown liquid        Grey solid
• Appearance in water:            -              Pale yellow          Orange-brown            Brown in KI
• Appearance in solvent:          -              Pale yellow          Orange-brown            Purple/pink

Disproportionation: similtaneous oxidation and reduction of halogen

Cold alkali

Equation: X2 + 2NaOH -> NaXO + NaX + H2O

Ionic equation: X2 + 2OH- -> XO- + X- + H2O 

0   +1    -1

e.g. I2 + 2NaOH -> NaIO (sodium iodate(I)) + NaI + H2O

Hot alkali

Equation: 3X2 + 6NaOH -> NaXO3 + 5NaX + 3H2O

Ionic equation: 3X2 + 6OH- -> XO3- + 5X- + 3H2O

0     +5        -1

e.g. 3Br2 + 6NaOH -> NaBrO3 (sodium bromate(III)) + 5NaBr + 3H2O

Metals: e.g. flourine and chlorine react with iron to form iron(III) halides

• Oxidation of iron: 2Fe -> 2Fe3+ + 6e-
• Reduction of chlorine: 3Cl2 + 6e- -> 6Cl-
• Overall: 2Fe(s) + 3Cl(g) -> 2FeCl3(s)
• Bromine is a weaker oxidising agent so a mixture of iron(II) and iron(III) bromide
• With iodine only iron(II) iodide is formed

Non-metals: e.g. chlorine reacts with sulfur to form sulfur(I) chloride

• S8(s) + 4Cl(g) -> 4S2Cl2(l)

Ions: all halogens except iodine oxidise Fe2+ tp Fe 3+ in solution

• 2Fe2+(aq) -> 2Fe3+(aq) + 2e-

Hydrogen:

• Chlorine - explodes in direct sunlight
• Bromine - 300C, platinum catalyst
• Iodine - 300C, platinum catalyst

Potassium halides with concentrated H2SO4: reducing power decreases down the group

• KF or KI: KF(s) + H2SO4(aq) -> KHSO4(s) + HF(g), (steamy fumes)
•        KCl(s) + H2SO4(aq) -> KHSO4(s) + HCl(g) (steamy white fumes)
  • reaction stops here because HCl/HF are not strong enough reducing agents. Not redox
• KBr: KBr(s) + H2SO4(aq) -> KHSO4(s) + HBr(g) (steamy fumes)
  • HBr is a stronger reducing agent than HCl so reacts with sulfuric acid in a redox reaction
•        2HBr(g) + H2SO4(aq) -> Br2(g) (orange fumes) + SO2(g) (choking fumes) + 2H2O(l)
• KI: KI(s) + H2SO4(aq) -> KHSO4(s) + HI(g)
•      2HI(g) + H2SO4(aq) -> I2(s) + SO2(g) + 2H2O(l) 
•      6HI + SO2(g) -> H2S(g) + 3I(s) + 2H2O(l) - HI reduces SO2 to H2S (toxic, egg smell)

Ammonia and water:

• Hydrogen halides are colourless gases.
• They are soluble and dissolve in water to make strong acids e.g.HCl(g) -> H+(aq) + Cl-(aq)
• They react with ammonia gas to produce white fumes e.g. NH3(g) + HCl(g) -> NH4Cl(s)

Displacement: more reactive halogen will displace less reactive halide ion below it

• Br2(aq) + 2KI(aq) -> 2KBr(aq) + I2(aq)

Test for halides in solution:

• Add dilute HNO3 - removes ions that could interfere with precipitate
• Add AgNO3 solution
• Precipitate of silver halide forms
• Test results with ammonia solution

Colours of silver halide precipitates:

• Chloride Cl-: white precipitate, dissolves in dilute NH3(aq), darkens in sunlight
• Bromide Br-: cream precipitate, dissolves in concentrated NH3, darkens in sunlight
• Iodide I-: yellow precipitate, insoluble in concentrated NH3, doesn't darken in sunlight

Decomposition in sunlight:

2AgBr(s) -> 2Ag(s) + Br2(l)

• Reaction is used in film photography - particles turn opaque silver upon exposure to light

Trends:

• Number of electrons increases down the group so van der Waal forces increase 
• Electronegativity decreases down the group 
• Oxidising ability decreases down the group 

Astatine

• Solid with highest boiling temperature
• Lowest electronegativity value
• Weakest oxidising agents

Fluorine

• Gas with lowest boiling temperature
• Highest electronegativity value
• Strongest oxidising agent