Topic 4 - Inorganic Flashcards
(40 cards)
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
