Inorganic Chemistry And The Periodic Table Flashcards
Group 2 + oxygen
General equation
Solid white oxide
2M(s) +O2(g)—-> 2MO(s)
Group 2 + chlorine
General equation
Solid white chlorides
M(s) + Cl2(g)——-> MCl2(s)
Group 2 + water
General equation
Metal hydroxide + hydrogen
M(s) + 2H2O(l) —-> M(OH)2 (aq) + H2(g)
Group 2 oxide + water
General equation
Hydroxides
MO(s) + H2O(l)—-> M(OH)2 (aq)
Group 2 oxide+ dilute acid
General equation
Salt+water
MO(s)+2HCl(aq)—-> MCl2(aq)+H2O(l)
Group 2 hydroxide + dilute acid
General equation
M(OH)2(aq)+2HCl(aq)—-> MCl2(aq)+2H2O(l)
Solubility trend group 2 hydroxides
Increases down the group
Solutions more strongly alkaline
Solubility trend group 2 sulfates
Decreases down the group
Trend in thermal stability down the group for G2 carbonates and nitrates
Increases
Polarising power of the cation decreases down the group as the ionic radius increases and so more shielding (lower charge density)
Less distortion of carbonate/ nitrate anion
Weakens the CO/ NO bond less
Thermal stability of G2 vsG1 nitrates/ carbonates
G2 carbonates/ nitrates are less thermally stable
The greater the charge on the cation, the greater the distortion and the less stable to anion becomes
2+ vs 1+
Thermal decomposition of G1 carbonates
Thermally stable (until higher temperatures)
Exception: Li2CO3 which decomposes to Li2O and CO2
Thermal decomposition of G2 carbonates
Form oxide and carbon dioxide
MCO3 (s)—-> MO(s) + CO2 (g)
Thermal decomposition of G1 nitrates
Forms nitrite and oxygen
2MNO3(s)——>2MNO2(s)+O2(g)
Exception:
LiNO3 which decomposes to form Li2O, NO2, O2
thermal decomposition of G2 nitrates
Forms oxide, nitrogen dioxide, oxygen
2M(NO2)2(s)——>2MO(s)+4NO2(g)+O2(g)
Testing the thermal stability of nitrates
How long it takes for a certain amount of oxygen to be produced (enough to relight a glowing splint)
How long it takes for an amount of brown gas (NO2) to be produced (in a fume cupboard)
Testing thermal stability of carbonates
How long it takes for an amount of carbon dioxide to be produced, use limewater
Limewater
Saturated solution of calcium hydroxide)
Formation of characteristic flame colours
Energy absorbed from the flams causes electrons to be promoted to higher energy levels, excitation. Electrons de excite and move back down to ground state, releasing energy in the form of light.
The difference in energy between the higher and lower energy levels determines the wavelength of light released and so the colour of light
Flame colours for:
Lithium, sodium, potassium, rubidium, caesium, calcium, strontium and barium
Red, orange/ yellow, lilac, red, blue, brick-red, crimson, green
States and colours of the halogens
F- pale yellow gas
Cl- green gas
Br- red/brown liquid
I- grey solid
Colour change when halogen displaces halide in aqueous solution
Cl displaces Br- yellow/ orange
Cl or Br displaces I- brown
Coloured layers of displaced halide after addition of an organic solvent
Bromine- orange/red
Iodine- pink/ violet
Chlorine is virtually colourless
Hypochlorous acid (HClO)
Ionises to make chlorate (1) ions
HClO+H2O——>ClO-+H3O+
ClO- kill bacteria
Disproportionation reaction of chlorine with cold, dilute aqueous sodium hydroxide to form bleach
Produces NaOX+NaX+H2O
Disproportionation reaction of chlorine with hot alkali
Produces NaXO3+5NaX+3H2O
Trend in reducing power of halides down the group
Increases
Because ions get bigger and there’s more shielding (outer electron more easily lost)
Reaction of KF or KCl with H2SO4
KCl(s)+H2SO4(l)—->KHSO4(s)+HCl(g)
Misty fumes
Reaction stops here
Not redox
Chlorine not strong enough reducing reagent to reduce the sulfur further
Reaction of KBr with H2SO4
KBr+H2SO4—>KHSO4+HBr misty fumes
2HBr(aq)+H2SO4(l)——>Br2(g)+SO2(g)+2H2O(l)
Redox
Sulfur oxidation state changes from +6 to +4
Br2 orange fumes
Reaction of KI with H2SO4
KI+H2SO4—>KHSO4+HI
2HI(g)+H2SO4—->I2(g)+SO2+2H2O
6HI(g)+SO2(g)—->H2S(g)+3I2(s)+@H2O(l)
+6 to +4 then +4 to -2
Reaction of hydrogen halides with ammonia and with water to produce acids
HX can dissolve in water to produce misty fumes of acidic gas (turn damp blue litmus paper red)
HX react with ammonia gas to produce white fumes
NH3(g)+HCl(g)——>NH4Cl(s)
reaction of aqueous Cl-, Br- and I- with aqueous silver nitrate solution, followed by aqueous ammonia solution
+dilute nitric acid to remove ions
+silver nitrate solution
AgF soluble so no precipitate
AgCl- white + ammonia solution- precipitate dissolves to give colourless solution
AgBr- cream + ammonia solution- unchanged with dilute, dissolves in concentrated to form colourless solution
AgI- yellow + ammonia solution- does not dissolve (precipitate remains)
Test for sulfates
+dilute HCl (removes carbonate ions)
+barium chloride solution
White precipitate of barium sulfate forms
Test for ammonium ions
+NaOH
Gently heat
Ammonia gas given off is alkaline so will dissolve in the water on damp red litmus paper and turn it blue
Equation for testing for ammonium ions
NH4Cl(aq)+NaOH(aq)—-> NH3(g)+H2O(l)+NaCl(aq)
Explain how the trend in reactivity of group 2 elements is determined by their electronic configurations
Outer electron further from nucleus
More shielding
So first ionisation energy decreases down group so reactivity increases
