Topic 4 - Inorganic Flashcards
Trend in ionisation energy down group 2
- atomic radius increases down the group due to additional electron shells
- increased shielding makes outer electrons easier to lose
- reactivity increases down the group
- ionisation energy decreases down the group
Trend in reactivity down group 2
- atomic radius increases down the group due to additional electron shells
- increased shielding makes outer electrons easier to lose
- low first and second ionisation energies
- reactivity increases down the group
Group 2 metal + water ->
metal hydroxide + hydrogen
M + 2H2O -> M(OH)2 + H2
Group 2 metal + oxygen ->
metal oxide
2M + O2 -> 2MO
Group 2 metal + chlorine ->
metal chloride
M + Cl2 -> MCl2
Group 2 metal oxide + water ->
metal hydroxide
MO + H2O -> M(OH)2
exceptions: beryllium and magnesium
Group 2 metal oxide + acid ->
salt and water
MO + 2HCl -> MCl2 + H2O
Group 2 metal hydroxide + acid ->
salt and water
M(OH)2 + 2HCl -> MCl2 + 2H2O
Solubility trend of group 2 hydroxides
generally, compounds of group 2 elements that contain singly charged negative ions (e.g. OH-) increase in solubility down the group
Solubility trend of group 2 sulfates
compounds that contain doubly charged negative ions (e.g. SO42- and CO32-) decrease in solubility down the group
exception: barium sulfate is insoluble
Thermal stability trend for groups 1 and 2 nitrates and carbonates
- carbonate and nitrate ions are large anions that can be made unstable by cations
- the cation polarises the anion, distorting it
- the greater the distortion, the less stable the compound
- large cations cause less distortion as they have a lower charge density
- the further down the group, the larger the cation, the lower the charge density, the lower the distortion and the more stable the compound
Thermal stability of group 1 vs group 2 compounds
- group 2 compounds are less thermally stable than group 1 compounds
- the greater the charge on the cation, the greater the distortion and the less stable the compound becomes
Decomposition of group 1 carbonates
thermally stable so cannot decompose
exception: Li2CO3, which decomposes to Li2O and CO2
Decomposition of group 1 nitrates
decompose to from the nitrite and oxygen (2MNO3 -> 2MNO2 + O2)
exception: LiNO3 to form Li2O, NO2 and O2
Decomposition of group 2 carbonates
decompose to form oxide and carbon dioxide
MCO3 -> MO + CO2
Decomposition of group 2 nitrates
decompose to form the oxide, nitrogen dioxide and oxygen
2M(NO3)2 -> 2MO + 4NO2 + O2
Flame colours of group 1 and 2 metals and their compounds
Li - red
Na - orange/yellow
K - lilac
Rb - red
Cs - blue
Ca - brick red
Sr - crimson
Ba - green
Flame colours explanation
- the energy absorbed from the flame cause electrons to move to higher energy levels
- the colours are seen as the electrons fall back down to lower energy levels, releasing energy in the form of light
- the difference in energy between the higher and lower levels determines the wavelength of the light released, which determines the colour of the light
Nitrate decomposition testing
- how long it takes until enough oxygen is produced to relight a glowing splint
- how long it take for brown gas (NO2) to form
Carbonate decomposition testing
how long it takes for carbon dioxide to be produced - limewater test
Flame test procedure
- mix a small amount of a compound you’re testing with a few drops of concentrated hydrochloric acid
- heat the nichrome wire in a hot Bunsen flame to clean it
- dip the wire into the compound/acid mixture and hold in a very hot flame
- note the colour produced
Group 7 electronegativity trend
- decreases down the group
- atomic radius increases down the group
- shielding effect increases
- nucleus has a weaker pull on shared electrons, leading to lower electronegativity
Group 7 melting and boiling temperature trend
- increase down the group
- increase in electron shells
- London forces between halogen molecules get stronger
- harder to overcome intermolecular forces
Group 7 reactivity trend
- decrease down the group
- group 7 atoms react by gaining an electron
- atomic radius increases down the group, so valence electrons are further from the nucleus
- increased shielding means nucleus cannot attract electrons as effectively
Trend in reactivity of Group 7 elements in terms of the redox reactions followed by the addition of an organic solvent
Oxidising power decreases moving down the group as ability to attract electrons decreases:
- Chlorine can displace bromide and iodide
- Bromine can displace iodide
- Iodide cannot displace either
E.g.
Cl₂ + 2Br- –> Br₂ + 2Cl-
Addition of cyclohexane:
- yellow-orange for bromine
- purple for iodine
- light green/yellow for chlorine
Disproportionation reaction of chlorine with water
Chlorine reacts with cold water:
produces chlorate (I) ions (ClO-) and chloride ions
chlorine is both oxidised and reduced, therefore disproportionation reaction
oxidisation state goes from 0 to both -1 and +1
Cl2 + H2O –> ClO- + Cl- + 2H+
Cl2 + H2O -> HCl + HClO
Use of chlorine in water treatment
The disproportionation reaction of water and chlorine is used to kill bacteria
This can pose risks as chlorine is toxic
disproportionation reaction of chlorine with cold, dilute aqueous sodium hydroxide to form bleach
forms sodium chloride and sodium chlorate(I)
2NaOH + Cl2 –> NaClO + NaCl + H2O
disproportionation reaction of chlorine with hot alkali
forms sodium chloride and sodium chlorate(v)
forms a species with oxidisation number -1 and +5 (chlorine)
3Cl2 + 6NaOH –> NaClO3 + 5NaCl + 3H2O
Trend in reducing power of halides down the group
reducing power increases down the group
- reducing agent gets oxidised (loses electron)
- attraction between halide’s nucleus and outer electrons gets weaker down the group due to shielding and increase in ionic radius
Reactions of KF or KCl with H2SO4
KF + H2SO4 -> KHSO4 + HF
KCl + H2SO4 -> KHSO4 + HCl
- misty fumes as hydrogen halide comes in contact with moisture in the air
- F and Cl aren’t strong enough to reduce sulfuric acid
- not a redox reaction
Reaction of KBr with H2SO4
KBr + H2SO4 -> KHSO4 + HBr
2HBr + H2SO4 -> Br2 + SO2 + 2H2O
- first reaction gives misty fumes of HBr gas
- Br- are a stronger reducing agent, and so react with H2SO4 in a redox reaction
Reaction of KI with H2SO4
KI + H2SO4 -> KHSO4 + HI
2HI + H2SO4 -> I2 + SO2 + 2H2O
6HI + SO2 -> H2S + 3I2 + 2H2O
- I- is a very good reducing agent and reduces SO2 to H2S
Precipitation reactions of halide ions with silver nitrate solution
- add dilute nitric acid to remove interfering ions
- add silver nitrate solution, and a precipitate is formed
Ag+ + X- -> AgX
fluoride - no precipitate as AgF is soluble
chloride - white
bromide - cream
iodide - yellow
addition of aqueous ammonia solution to silver halide precipitates
ammonia used to tell the halides apart
AgCl - dissolves to form colourless solution
AgBr - remains unchanged in dilute solution, dissolves in concentrated solution
AgI - does not dissolve
reaction of hydrogen halide with water
- dissolve in water to produce misty fumes of acidic gas (turns damp blue litmus paper red)
- hydrochloric, hydrobromic, or hydroiodic acid
HCl -> H+ + Cl-
HCl + H2O -> H3O+ + Cl-
reaction of hydrogen halide with ammonia
- gives white fumes
e.g. hydrogen chloride gives ammonium chloride
NH3 + HCl -> NH4Cl
Test for carbonate ions using hydrochloric acid
- CO32- and HCO3-
- both fizz to give off carbon dioxide
CO32- + 2H+ -> CO2 + H2O
HCO3- + H+ -> CO2 + H2O
- test for CO2 using limewater (turns cloudy)
Test for sulfates with hydrochloric acid and barium chloride
- add dilute HCl followed by barium chloride solution (BaCl2)
- HCl is used to remove traces of carbonate ions
Ba2+ + SO42- -> BaSO4
- white precipitate of barium sulfate = original compound contained a sulfate
Test for ammonia using litmus paper and NaOH
- damp red litmus paper turns blue
- add sodium hydroxide to the substance and gently heat the mixture
- if ammonia is given off, the substance contains ammonium ions
NH4+ + OH- -> NH3 + H2O
e.g. NH4Cl + NaOH -> NH3 + H2O + NaCl
