GROUP II ELEMENTS

period	element	Atomic #	Elec config
2	Beryllium, Be	4	1s22s2
3	Magnesium, Mg	12	1s22s22p63s2
4	Calcium, Ca	20	1s22s22p63s23p64s2
5	Strontium, Sr	38	[Kr] 5s2
6	Barium, Ba	56	[Xe] 6s2
7	Radium, Ra	88	[Ra] 7s2

Atomic & Ionic radius


-increase down group II
-due to increased screening effect which also cancels out increased nuclear charge attraction > more repulsion > larger radius


Melting pt

mp of Mg > Ba: less than 850°C much less than transition metals (mp: >1000°C) due to large atomic radii > valence s-electrons less tightly attracted > weaker metallic bonds > lower mp
mp of Ca > Ba: decrease due to increased atomic radii > weaker metallic bonds

Oxidation #
all group II elements: oxid # +2 due to same valence shell config (n s²) > M²⁺ ions formed in reactions: M(s) - 2e^- > M²⁺(aq)

1st Ionisation Energy

1st IE decreases down group due to increase in screening effect of inner filled electron shells which is greater than increased nuclear charge attraction > larger atomic radius > valence e^- further from nucleus > less attraction > less energy needed to remove it

Group II elements as reducing agents
group II elements: v.electropositive & reactive as reducing agents
reactivity increases w/ atomic # (Mg < Ca < Sr < Ba)
reactivity depends on IE: down the group atomic radius increases > decreasing IE > valence electrons more easily lost > ion formed more readily > stronger reducing agent
reactivity also explained in terms of standard potential,
more -ve standard electrode potential > more reactive reducing agent
being reducing agents: M - 2e^- >(oxidised)> M²⁺(aq)

Reactivity w/ water

	reaction w/ water	product
Mg	v.slow w/ water at room temp, more readily w/ steam Mg(s) + H2O(g) >(heat)> MgO(s) + H2(g)	MgO- basic oxide, insoluble in water
Ca	readily w/ cold water: Ca + H2O > Ca(OH)2(s) + H2(g)	Ca(OH)2- basic hydroxide, slightly soluble in H2O
Sr	readily w/ cold water: Sr + H2O > Sr(OH)2(aq) + H2(g)	Sr(OH)2- basic hydroxide, soluble in water
Ba	violently w/ cold water: Ba + H2O > Ba(OH)2(aq) + H2(g)	Ba(OH)2- basic hydroxide, soluble in water

reactivity w/ water: Mg < Ca < Sr < Ba
basic strength of aq hydroxide: Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂

Reactivity w/ oxygen
all form white solid oxides when heated in oxygen: 2M(s) + O₂(g) > 2MO(s)
reactivity increases down group: Mg < Ca < Sr < Ba
w/ excess O₂: Sr & Ba react > oxides + peroxides: Sr(s) + O₂(g)(excess) > SrO₂, Ba(s) + O₂(g)(excess) > BaO₂
thin layer of oxides cover each group II metal due to oxygen in air

Reactivity of group II oxides w/ water
all oxides react w/ water > hydroxides
MgO(s) + H₂O(l) > Mg(OH)₂(s)
CaO(s) + H₂O(l) > Ca(OH)₂(s)
SrO(s) + H₂O(l) > Sr(OH)₂(aq)
BaO(s) + H₂O(l) > Ba(OH)₂(aq)

	reaction w/ water	product
MgO	MgO(s) + H2O(l) > Mg(OH)2(s) Mg(OH)2(s) Mg2+(aq) + 2OH-(aq)(low conc)	MgO: sparingly soluble > weakly basic, pH ~8
CaO	vigorous + exothermic: CaO(s) + H2O(l) > Ca(OH)2(s) Ca(OH)2(s) + aq Ca(OH)2(aq) Ca2+(aq) + 2OH-(aq)	Ca(OH)2: slightly soluble > weakly alkaline sol called limewater
SrO	vigorous + exothermic: SrO(s) + H2O(l) > Sr(OH)2(aq) Sr(OH)2(s) + aq Sr(OH)2(aq) Sr2+(aq) + 2OH-(aq)	Sr(OH)2: v.soluble > strong alkaline, pH ~13
BaO	vigorous + exothermic: BaO(s) + H2O(l) > Ba(OH)2(aq) Ba(OH)2(s) + aq Ba(OH)2(aq) Ba2+(aq) + 2OH-(aq)	Ba(OH)2: v.soluble >strong alkaline, pH ~13

CaO(s) = quicklime
CaO(s) + water > swells and cracks > v.hot > crumbles to fine white bulky powder = slaked lime
BaO + water > a lot of heat > becomes incandescent (it glows) > excess water > steam

Thermal decomposition of group II carbonates

carbonates >(heat > decompose)> metal oxides + CO₂
MgCO₃(s) >(heat)> MgO(s) + CO₂(g)
CaCO₃(s) >(heat)> CaO(s) + CO₂(g)
SrCO₃(s) >(heat)> SrO(s) + CO₂(g)
BaCO₃(s) >(heat)> BaO(s) + CO₂(g)
thermal stability increases down group [MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃ (most stable)] due to less distortion of carbonate ion, CO₃^2- > oxide less easily formed (down group increasing atomic radius > smaller charge density > less ability to polarise and distort carbonate ion)


Thermal decomposition of nitrates
nitrate >(heat)> oxide + NO₂ + O₂
2Mg(NO₃)₂ >(heat)> 2MgO + 4NO₂ + O₂
2Ca(NO₃)₂ >(heat)> 2CaO + 4NO₂ + O₂
2Sr(NO₃)₂ >(heat)> 2SrO + 4NO₂ + O₂
2Ba(NO₃)₂ >(heat)> 2BaO + 4NO₂ + O₂
thermal stability increases down the group due to less distortion of nitrate ion, NO₃^- > oxide less easily formed (down group increasing atomic radius > smaller charge density > less ability to polarise and distort nitrate ion)

Solubility of sulphates
solubility depends on:
-lattice energy: energy absorbed to separate the 2 ions (anion & cation), larger lattice energy > ions held stongly together > lower solubility

-hydration energy of cations: lower hydration energy > less soluble

down group II solubility decreases:
  -lattice energy: ~ const down group II > doesn't affect solubiliuty v.much
  -hydration energy of cations: decreases down group II due to larger ionic radius > less energy released when cation hydrolysed


Uses of group II cpds
MgO: v. high mp (2900°C) > refractory in lining furnaces
Mg(OH)₂: in milk of magnesia as antacid and laxative
CaO: neutralises acidic soil
CaCO₃: marble as building material
BaSO₄: in 'barium meal; for gastrointestineal x-ray photography as it is opaque to x-ray [Ba²⁺ is poisonous, but since BaSO₄ is practically insoluble > minute amts in body > not enough to harm]


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