TOPIC : GROUP 2

1. Group 2 elements are known as alkaline earth metals or s-block elements as their valence electrons are in s orbitals.2. The Group 2 elements are beryllium, magnesium, calcium, strontium, barium and 3. Some information of the Group 2 elements is listed in the table below.

Element Be Mg Ca Sr BaName Beryllium Magnesium Calcium Strontium BariumProton number

4 12 20 38 56

Electronic configuration

2.2[He] 2s2

2.8.2[Ne] 3s2

28.8.2[Ar] 4s2

2.8.18.8.2[Kr] 5s2

2.8.18.18.8.2[Xe] 6s2

4. All have valence shell electronic configuration of ns2

5. They are all reactive metals and are not found in the free elemental states in nature.6. In their pure state, they have a silver colour but tarnish rapidly in air due to the formation of an oxide layer on the metals’ surface.

Example : 2Mg(s) + O2(g) 2MgO(s)

7. They are soft and can be easily cut with a knife.8. The Group 2 elements give characteristic flame test :

Element Magnesium Calcium Strontium BariumColour of the flame

Brilliant white Brick red Crimson Apple green

Physical Properties1. Atomic radius increase down the group as the outermost electrons are in the shells further from the nucleus and the screening effect of inner electrons increases.2. Ionic radius (of M2+) increases down the group (same reason as for increase in atomic radius). Ionic radius is less then the original atom as the nuclear charge exceeds the electronic charge.3. Melting points generally decrease down the group as melting bonding gets weaker due to increased atomic size (the number of delocalized electrons per atom is the same). However, different crystalline structures can also affect the melting point. (Note : The melting point of magnesium is lower than that of calcium).4. Ionisation energies (1st and 2nd) are relatively low and decreases down the group (due to increasing size and increasing shielding effect of inner electrons).5. Electronegativity increases down the group as increased size and increased screening effect make any shared pair of electrons less strongly attracted to the nucleus.Chemical Properties1. General(a) The group 2 elements are typical metals and are very reactive, powerful reducing agents.

M M2+ + 2e(b) They form mostly ionic compounds containing M2+ ions except for some covalent compounds formed by beryllium; and only have the oxidation number +2 in all their compounds.(c) On going down the group, the elements generally become more reactive because as proton number increases, the atomic radius increases, the ionization energy decreses, and atom loses its electrons more readily to form the M2+ ion. So the metals become more electropositive.

2. Reaction With Water (a) The reactivity of the elements with water increases down the group. As the elements become more electropositive their reducing power increases; reducing water to hydrogen an forming their hydroxides or oxides. (b) (i) Beryllium does not react with cold water. (will react slowly with steam at very high temperature)

Be(s) + H2O(l) BeO(s) + H2(g)

(ii) Heated magnesium reacts rapidly with steam to form magnesium oxide and hydrogen.

Mg(s) + H2O(l) MgO(s) + H2(g)

(iii) Calcium reacts very slowly in cold water but rapidly with hot water to form the sparingly soluble calcium hydroxide and hydrogen.

Ca(s) + 2H2O(l) Ca(OH)2(s) + H2(g)

(iv) Strontium and barium react vigorously with cold water to form their respective hydroxide solution and hydrogen.

Sr(s) + 2H2O(l) Sr(OH)2(s) + H2(g)

Ba(s) + 2H2O(l) Ba(OH)2(s) + H2(g)

3. Reaction with oxygen (a) The metal (when heated) burn in air to form the oxides which are white solids.

General equation : 2M(s) + O2(g) 2MO(s)

(b) Reactivity increases down the group. Fine barium powder will burn spontaneously in air, (Barium is normally stored in paraffin oil).

2Be(s) + O2(g) 2BeO(s)

2Mg(s) + O2(g) 2MgO(s)

2Ca(s) + O2(g) 2CaO(s)

2Sr(s) + O2(g) 2SrO(s)

2Ba(s) + O2(g) 2BaO(s)

(c) Due to the decreasing polarizing power of the cation down the group, the formation of peroxides becomes possible. Strontium and barium from SrO2 and BaO2 respectively with excess oxygen but magnesium does not.

Sr(s) + O2(g) SrO2(s)

Ba(s) + O2(g) BaO2(s)

(d) Properties of the oxides (i) the oxides are all basic (except BeO which is amphoteric)

(ii) the solubility of the oxides in water increases down the group : BeO is insoluble MgO is very sparingly soluble CaO is sparingly soluble SrO soluble BaO soluble Forming the alkaline hydroxide Solubility of oxides depends on solubility of hydroxides that increases down the group.

MgO(s) + H2O(l) Mg(OH)2(aq)

(sparingly soluble. Weakly alkaline)

CaO(s) + H2O(l) Ca(OH)2(aq)

SrO(s) + H2O(l) Sr(OH)2(aq)

BaO(s) + H2O(l) Ba(OH)2(aq)

(strongly alkaline)

Trend in thermal stability of the Nitrates, Carbonates and Hydroxides of Group 2 elements

All nitrates, carbonates and hydroxides of the Group 2 elements decompose when heated strongly. Thermal stability increases down the group, that is ease of thermal decomposition decreases.

1. Nitrates undergo thermal decomposition to form the corresponding oxide, nitrogen dioxide gas and oxygen.

NO2 -- brown gasO2 -- ignites a glowing splinter

2M(NO3)2(s)

Trend in thermal stability of the nitrates, carbonates and hydroxides of Group 2 elementsExplanation :

1. (a) Thermal decomposition occurs due to the polarization of the large anion by the Group 2 cation of high charge density. (b) Polarisation distorts the electron cloud and weakens bonds, resulting in decomposition. For example, the large carbonate ion decomposes as shown below

Polarisation occurs when a small cation distorts the electron cloud of a large anion. The polarizing effect is greastest between a small cation with a high charge (high charge density) and a large anion.

2. (a) The thermal stability of Group 2 compounds increases down the group from Be to Ba. On descending the group, temperatures required to decompose the compounds increase.

(b) Reason : (i) the size of Group 2 cations increases down the group. As all the ions have 2+ charges, the charge density and hence the polarizing power of the cations also decreases down the group. (ii) When the polarization of large anions decreases, they are not as easy decomposed. (c) The compounds of beryllium are the least stable and undergo the thermal decomposition most easily.

Mg 2+

2+

Ca

2+

Ba

3. (a) When the nitrates, carbonates and hydroxides of the elements are heated, the metal oxides are formed because they are more stable. (b) Reasons for extra stability of the metal oxide: (i) The oxide ion, O2-, is smaller and not as easily polarized compared to the bigger NO3

- , CO32- and OH- ions.

(ii) The ionic bond between the cation (M2+) and the smaller oxide ion will be stronger resulting in a higher lattice energy for the metal oxide.

Solubility Of Sulphates1. Sulphates of Group 2 elements are formed when the basic oxides react with sulphuric acid. The sulphate has the general formula MSO4 (e.g. MgSO4 , CaSO4)2. The solubility of the sulphates decreases down the group. Magnesium sulphate is very soluble in water, but barium sulphate is very sparingly soluble (insoluble).3. The solubility of an ionic compound depends on two factors : (a) The lattice energy of the ionic solid (b) The enthalpy of hydration (or hydration energy) of the carbon anion.4. (a) A higher lattice energy will result in a decrease in solubility. (i) Energy is required to break the solid lattice apart into its free ions.

M2+X2-(s) M2+

(g) + X2-(g)

This process is endothermic and the amount of energy required per mole of compound is equivalent to its lattice energy.

(Lattice energy is defined as the energy evolved (exothermic) when 1 mol of an ionic lattice is formed from its constituent gaseous ions). (ii) Lattice energy increases when the size of the ion(s) decreases, the charge on the ion(s) increases; and its charge density increases as more energy is required to break the stronger ionic bonds that hold the ions together.

(b) A higher enthalpy will result in an increase in solubility (i) When an ionic compound dissolves, the free ions become hydrated by water molecules, heat energy is given out. the energy involved is the enthalpy of hydration.

M2+(g) + aq M2+

(aq) ; X2-(g) + aq X2-

(aq) (ii) Enthalpy of hydration increases when the size of the ion decreases, the charge on the ion increases; and its charge density increases.

5. (a) Going down the group, the size of the M2+ ion increases and charge density decreases. Hence both the lattice energy and the enthalpy oh hydration decrease. (b) However, for the sulphates of the Group 2 elements, the decrease in lattice energy down the group is very small. This is because the size of the sulphate ion is very large compared to the size of the cation M2+.

So the decrease in lattice energy of the sulphate down the group is small and will not contribute much to its solubility. (c) However the increase in ionic size of the cations from Be2+ to Ba2+ causes a large decrease in the enthalpy of hydration of the cations down the group. A decrease in enthalpy of hydration causes the solubility of the sulphates to decrease. (d) In summary, going down Group 2, the enthalpy of the hydration of the cations on the lattice energy decrease because the size of the cations increases and the charge density decreases. However, the enthalpy of hydration decreases much more than the lattice energy. Hence the solubility of the sulphates

decreases down the group.Anomalous Properties of Beryllium1. Some of the properties of beryllium are more similar to the properties of aluminium (Group 13) tha those of the elements of Group 2. These properties atypical of the Group 2 elements arise due to

(a) the small atomic radius of Be resulting in its very high ionization energy.(b) the high charge density of the Be2+ ion giving it a high polarizing power. As a

result, beryllium forms covalent compounds as well as ionic compounds with covalent character.

(c) the high charge density of the Be2+ ion also enables it to attract lone pair electrons from ligands to form complex ion.(d) the electronegativity of be is equal to that of aluminium.

2. Beryllium chloride is a covalent solid with low melting point whereas all the other chlorides of Group 2 elements are ionic compounds. At room temperature beryllium chlorides exists as Be2Cl4

4-

It Hydrolyses in water forming fumes of HCl

3. Beryllium oxides is amphoteric while all the other oxides of Group 2 elements are basic. 4. Beryllium forms complex ions such as [BeF4]2- and [Be(OH)4]2- but the elements of Group 2 do not.

5. Beryllium metal reacts with hot concentrated sodium hydroxide to form salt sodium beryllate and hydrogen.