AS Chemistry - Redox Reactions and Group 2 Elements

Author:
Emily
Updated date:
Mar 24, 2020

Emily is a PhD student with a BSc in Molecular Biology and Genetics, an MSc in Molecular Medicine, and a passion for science communication.

Redox Reactions

A redox reaction is a reaction that involves both oxidation (the loss of electrons) and reduction (the gain of electrons).

In order to identify whether a reaction is redox or not, you can write separate half equations that show how electrons are lost/gained.
For example, take the equation for the reaction of Calcium and Oxygen:

2Ca (g) + O₂(g) -> 2CaO (s),

The two half equations for this reaction are:

Ca -> Ca ²⁺ + 2e^-(this reaction shows oxidation).
and
O₂+ 4e^--> 2O^2-(this reaction shows reduction)

Therefore, you can conclude that this reaction is a redox reaction because it involves both reduction and oxidation.

There is another way of identifying a redox reaction (I personally find this method easier), in which you apply oxidation numbers to the equation to work out what has been oxidised and what has been reduced.

Reaction of Calcium and Oxygen

Basically there are 10 rules that show which elements and their oxidation numbers take priority in a reaction.
So, in order of importance, the 10 rules are as follows:

1) Group 1 elements (all have an oxidation number of +1)

2) Group 2 elements (all have an oxidation number of +2)

3) Group 3 elements (all have an oxidation number of +3)

4) Flourine (with an oxidation number of -1)

5) Hydrogen (with an oxidation number of +1)

6) Oxygen (with an oxidation number of -2)

7) Chlorine (with an oxidation number of -1)

8) Group 7 elements (all have an oxidation number of -1), Group 6 elements (-2) and Group 5 elements (-3).

Element	First Ionisation Energy/kJ mol-1
Be	900
Mg	736
Ca	590
Sr	548
Ba	502
Ra	509

Group 2

Also known as the alkaline earth metals, group 2 consist of the elements Beryllium, Magnesium, Calcium, Strontium and Barium.
They all have reasonably high melting and boiling points, low densities and they all form colourless compounds.
Together with group 1 (the alkali metals), they form the s block of the periodic table because their highest energy electrons are all in s sub-shells (a spherical orbital capable of holding 2 electrons). This means that the alkaline earth metals have 2 electrons in their outer shells.

Trends

Reactivity increases down group 2, this is due to 3 things:

1) The electron shielding increases as you go down the group.

2) The atomic radii also increases.

3) Nuclear charge increases (because of the increasing number of protons), however this is overpowered by the nuclear charge and atomic radii.

Basically, the more electron shielding an atom has the less attracted it's outermost electrons are to the positive nucleus and thus the electrons are lost easier.

From this we can deduce that the ionisation energy decreases as we go down the group.

Above is a table to show the ionisation energies of the group 2 elements.

Water and Group 2 Elements

Reactions

Oxygen:

All of the elements in group 2 react vigorously with Oxygen, the product of which is an ionic oxide. The general formula for this reaction is MO (where M is the group 2 element).
For example, Magnesium reacts with Oxygen to form Magnesium Oxide the formula for which is:

2Mg (s) + O₂(g) 2MgO (s)

This is a redox reaction.

Water:

All of the group 2 elements form hydroxides when reacted with water. The general formula for these reactions is M(OH)₂ (where M is the group 2 element). Hydrogen is given off during these reactions.
For example, Magnesium reacts with water to form Magnesium Hydroxide and Hydrogen gas in the following equation:

Mg (s) + 2H₂O(g) -> Mg(OH)₂(aq) + H₂(g)

This is also a redox reaction.