The elements in Group 2
- are metals (beryllium is a semi-metal 0r metalloid)
- have good conductivity of electricity (have delocalised electrons)
- have high melting and boiling points
- have silver-grey metallic appearance
Note: Radium is too radioactive to handle in a normal laboratory. Its other properties closely resemble those of barium.
They form ionic compounds* which
- are colourless in solution (white when solid)
- have high melting and boiling points
- conduct electricity when molten or in solution (when the ions are mobile and able to act as charge carriers)
*… except for beryllium which behaves differently because of its very small size, forming only covalent compounds.
To understand the reactions of Group 2 it helps to consider their electron arrangements:
Be [He] 2s2 1s2 2s2
Mg [Ne] 3s2 1s2 2s2 2p6 3s2
Ca [Ar] 4s2 1s2 2s2 2p6 3s2 3p6 4s2
Sr [Kr] 5s2 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2
Ba [Xe] 6s2 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s2
All the atoms of Group 2 elements have two electrons in their outer shell. Group 2 elements react by losing two electrons to form M2+ ions:
M(g) → M+(g) + e– ΔHI1
M+(g) → M2+(g) + e– ΔHI2
The energy required to remove the first of these electrons is the first ionisation energy, and the energy required to remove the second electron is the second ionisation energy. Note that ionisation energies are defined for moles of atoms in the gaseous state, so this must be reflected in the state symbols:
It is the sum of these ionisation energies which determines the relative reactivity of the elements:
|Element||ΔHI1 kJmol-1||ΔHI2 kJmol-1||Total kJmol-1|
It is clear from the data that we would expect reactivity to increase down the group as the energy required to remove the two outer shell electrons decreases and the 2+ ion becomes easier to form. We can explain this in terms of atomic radius, shielding and nuclear charge:
Down the group:
- the atomic radius increases with increasing number of shells, so the outer electrons are further from the attraction of the nucleus
- there are more filled inner shells shielding the two outer shell electrons from the nuclear charge
- these factors outweigh the increasing nuclear charge due to the increasing number of protons in the nucleus (atomic number)
- thus the energy required to form the M2+ ion decreases
Reactions of Group 2 elements
Experimental evidence for this order of reactivity comes from observing the redox reactions of the Group 2 metals with oxygen, water and dilute acids.
1) with oxygen
Group 2 metals undergo redox reactions with oxygen from the air (the most reactive of them are kept under oil like the Group 1 metals) and will burn with characteristic flame colours. [Mg = white, Ca = brick red, Sr = crimson, Ba = apple green]
2) with water
Group 2 metals Ca, Sr and Ba undergo redox reactions readily with water to form hydroxides, and give off hydrogen gas. These reactions are exothermic.
Mg reacts very slowly with water (over several days), producing a white suspension of magnesium hydroxide rather than a solution, because magnesium hydroxide is barely soluble in water.
3) with dilute acids
The Group 2 metals undergo redox reaction with dilute acids to form salts and hydrogen gas. The reactions become more vigorous and exothermic down the group.
The Group 2 elements are reducing agents. It is clear in each of their reactions above that they have themselves been oxidised, causing the substance they have reacted with to be reduced. The more reactive they are, the better they are at doing this, so Group 2 become better reducing agents, having greater reducing power as we go down the group.
Reactions of Group 2 Oxides
The oxides of Group 2 metals react with water to form metal hydroxides. These are NOT redox reactions.
CaO(s) + H2O(l) → Ca2+(aq) + 2OH–(aq)
Note that these hydroxides are only sparingly soluble, so once the solution has become saturated, the metal hydroxide forms as a precipitate.
Ca2+(aq) + 2OH–(aq) → Ca(OH)2(s)
The presence of hydroxide ions in the aqueous solution makes the resulting solutions alkaline, with pH values in the range 9 – 13.
Going down the group the metal hydroxides become more soluble in water, so the concentration of hydroxide ions that can be dissolved increases down the group, making the solution more alkaline and their pH higher.
Uses of Group 2 compounds
Calcium hydroxide (slaked lime) is used in agriculture to increase the pH of acidic soils:
Ca(OH)2(s) + 2H+(aq) → Ca2+(aq) + 2H2O(l)
Group 2 hydroxides and carbonates are bases, and are often used as antacids to treat acid indigestion (an excess of hydrochloric acid in the stomach) by neutralisation. Gaviscon and Rennies use calcium carbonate and magnesium carbonate, while Milk of Magnesia is a suspension of magnesium hydroxide in water.
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + 2H2O(l)