3.1.2(a)-(c) - Group 2 reactivity and ionisation energy

From s² electrons to 2+ ions

The Group 2 elements in this lesson are Mg, Ca, Sr and Ba. Each atom has two electrons in its outer s sub-shell, so the outer-shell pattern is s².

In their redox reactions, these metals lose the two outer s electrons to form 2+ ions:

The metal has lost electrons, so the metal has been oxidised. The species that gains those electrons has been reduced. This is why the reactions in this lesson are redox reactions.

Worked example: magnesium forming a 2+ ion

Magnesium has outer-shell configuration 3s². In a redox reaction it loses both outer electrons:

The oxidation number of Mg increases from 0 in the element to +2 in Mg²⁺, so magnesium is oxidised.

Notice what is not happening: magnesium does not usually form Mg⁺ in these reactions. Losing one electron would leave one electron still in the 3s sub-shell; the common Group 2 ion is Mg²⁺.

Reactions with oxygen

For the oxide-forming reactions specified here, the product has formula MO because the metal ion is M²⁺ and the oxide ion is O²⁻. These equations model formation of the monoxide; if a question explicitly supplies a different oxygen-containing product, use its supplied formula rather than assuming that every set of conditions must give MO.

The general equation is:

Examples:

This is a redox reaction. The metal changes from oxidation number 0 in the element to +2 in the oxide. Oxygen changes from 0 in O₂ to -2 in the oxide ion. The metal is oxidised and oxygen is reduced.

Worked example: barium with oxygen

Barium forms Ba²⁺ and oxygen forms O²⁻, so the product formula is BaO for the specified oxide-forming reaction. Start with:

Balance oxygen atoms by making two barium oxide formula units:

Then balance barium:

Magnesium burns with an intense white light and leaves white magnesium oxide. A white solid is an observation; “magnesium is oxidised” is the electron-transfer inference. For the oxide-forming equations here use $\ce{MO}$; do not infer a product formula solely from the brightness of the flame.

The relative reactivity increases from magnesium to barium. In practice, the reactions become easier to start and more vigorous down the group.

Reactions with water and dilute acids

With cold water, calcium, strontium and barium react to form a metal hydroxide and hydrogen. Magnesium reacts only very slowly with cold water; if steam is specified, magnesium can form magnesium oxide and hydrogen instead. Always use the condition given in the question.

For the cold-water hydroxide pattern:

Examples:

The formula M(OH)₂ matters. A Group 2 metal ion has charge 2+, while each hydroxide ion is OH⁻, so two hydroxide ions are needed. Writing BaOH is a common formula error.

The reaction with water is redox:

• the metal is oxidised from 0 to +2
• hydrogen in water is reduced from +1 to 0 in H₂
• oxygen remains at -2

The steam reaction is $\ce{Mg(s) + H2O(g) -> MgO(s) + H2(g)}$. The product is oxide, whereas slow reaction with cold liquid water gives hydroxide. Effervescence alone shows gas formation; a small collected sample giving a squeaky pop with a lighted splint supports identification as hydrogen. Calcium can give a cloudy white mixture as slightly soluble $\ce{Ca(OH)2}$ forms; gas bubbles and a precipitate are different observations.

Dilute acid reactions

With a non-oxidising dilute acid such as hydrochloric acid, these metals follow a salt-plus-hydrogen pattern. This is not a claim that every acid or reaction condition behaves identically; use the acid and conditions actually specified.

With dilute hydrochloric acid:

Example:

The redox core is:

The metal loses two electrons and becomes M²⁺. Two hydrogen ions gain those electrons and form one molecule of hydrogen gas.

Worked example: calcium with dilute hydrochloric acid

Calcium forms Ca²⁺, so the chloride salt must be CaCl₂. Hydrogen gas is diatomic, so it is H₂.

Start with the products:

Two chloride ions are needed for CaCl₂, so place a 2 before HCl:

The same pattern becomes more vigorous down the group from Mg to Ba, shown by faster effervescence and faster disappearance of the metal when comparable pieces and acid conditions are used.