The Reactivity Series

Explaining Reactivity

Why are some metals so much more explosive than others? The fundamental principle is that when metals react, they lose their outer shell electrons to form positive ions (also known as cations). This loss of electrons is called oxidation.

A metal's reactivity is determined by how easily it loses these outer electrons. Metals at the top of the series lose their electrons much more readily than those at the bottom.

For alkali metals (Group 1 elements like Li, Na, K), reactivity increases as you go down the group. This is because the atomic radius gets larger, and more internal electron shells shield the nucleus.

The weaker electrostatic attraction between the positive nucleus and the negative outer electron means the outer electron is lost more easily. Group 1 metals are also more reactive than Group 2 metals (like Mg or Ca) because they only need to lose one electron instead of two.

Reactions of Metals with Water

You can safely wash a copper coin in the sink, but dropping potassium in water causes an explosion. When the most reactive metals (K, Na, Li, Ca) react with cold water, they produce a metal hydroxide and hydrogen gas. The presence of hydrogen gas is confirmed by testing with a lighted splint, which produces a "squeaky pop".

The general equation for Group 1 metals reacting with water is:

$$2M(s) + 2H_2O(l) \rightarrow 2MOH(aq) + H_2(g)$$

Visual observations vary depending on the metal's reactivity:

• Potassium: Reacts violently, floats and moves very rapidly, produces a lilac flame, melts into a ball, and fizzes vigorously.
• Sodium: Reacts vigorously, floats and moves rapidly on the surface, melts into a shiny ball, and fizzes.
• Lithium: Reacts steadily, floats and moves slowly, fizzes, and slowly disappears (does not usually melt).
• Calcium: Reacts moderately, sinks to the bottom, produces bubbles of gas, and the solution turns cloudy white because the calcium hydroxide forms a sparingly soluble precipitate.

Magnesium reacts extremely slowly with cold water. However, it and other less reactive metals can react with steam to produce a solid metal oxide and hydrogen:

• Magnesium: Reacts vigorously with steam, burning with a bright white flame to produce white magnesium oxide ($MgO$).
• Zinc: Reacts when heated strongly to produce zinc oxide ($ZnO$), which is yellow when hot and white when cold.
• Iron: Reacts slowly when "red hot" to produce black triiron tetraoxide ($Fe_3O_4$).
• Copper: Has absolutely no reaction with water or steam.

Reactions of Metals with Dilute Acids

If you drop a piece of metal into acid, the speed of the bubbles tells you exactly how reactive it is. When metals react with dilute hydrochloric acid ($HCl$) or dilute sulfuric acid ($H_2SO_4$), they produce a metal salt and hydrogen gas. Hydrochloric acid produces salts called chlorides, while sulfuric acid produces salts called sulfates.

The observations depend heavily on the metal's position in the reactivity series:

• Potassium and Sodium: React explosively and violently. These reactions are too dangerous to perform in school labs.
• Lithium: Reacts violently with rapid effervescence (bubbling).
• Calcium and Magnesium: React very vigorously with rapid effervescence. The metal disappears quickly, and the mixture gets very hot (highly exothermic).
• Zinc: Reacts steadily, producing continuous bubbles of hydrogen gas, and the metal disappears slowly.
• Iron: Reacts slowly with a slow production of bubbles. It forms an Iron(II) salt, such as iron(II) chloride ($FeCl_2$).
• Copper: Shows no reaction, produces no bubbles, and causes no temperature change.

Deducing Reactivity from Experiments

How do chemists actually figure out the reactivity series? They perform controlled experiments to compare the vigour of reactions. This is done by measuring the rate of hydrogen gas production (effervescence) or the temperature change ($\Delta T$) when metals are added to acids or salt solutions.

A greater temperature rise indicates a more reactive metal. To measure this accurately, scientists use calorimetry, often carrying out the reaction in a polystyrene cup. The polystyrene acts as a thermal insulator, reducing heat loss to the surroundings and giving a more accurate temperature reading.

When comparing metals, it is crucial to keep control variables the same to ensure a fair test. You must use the same volume and concentration of acid, the same mass and surface area of the metal, and the same initial temperature.

Interpreting Experimental Data To deduce the order of reactivity, you must analyse experimental results.

1. Analysing Temperature Changes A student adds different metals to dilute acid and measures the maximum temperature increase ($\Delta T$):

• Metal A: $14^\circ C$
• Metal B: $0^\circ C$
• Metal C: $46^\circ C$
• Metal D: $22^\circ C$

Analysis: The order of reactivity (most to least) is C > D > A > B. Metal C caused the largest temperature rise, meaning it is the most reactive. Metal B caused no temperature rise, meaning it did not react (e.g., copper).

2. Analysing Displacement Matrices Chemists use displacement tables to rank metals. If a metal is added to a salt solution and a reaction occurs, the metal is more reactive than the metal in the salt.

Metal	Salt Solution of A	Salt Solution of B	Salt Solution of C
Metal A	No reaction	Reaction	Reaction
Metal B	No reaction	No reaction	No reaction
Metal C	No reaction	Reaction	No reaction

Analysis:

• Metal A displaces B and C, so it is the most reactive.
• Metal C displaces B but not A, so it is in the middle.
• Metal B cannot displace any metal, so it is the least reactive. Order of reactivity: A > C > B.

Displacement Reactions

In chemistry, a more reactive element will always bully a less reactive one out of a compound. A displacement reaction occurs when a more reactive metal takes the place of a less reactive metal in an aqueous salt solution or a solid metal oxide.

When a solid metal is added to an aqueous salt solution of a less reactive metal, distinct visual changes occur. For example, when magnesium is added to blue copper(II) sulfate solution:

• The blue solution fades to colourless as magnesium sulfate forms.
• A pink-brown solid forms, which is the displaced copper.
• The piece of magnesium appears to get smaller or disappear.

The ionic equation for this reaction omits spectator ions (like the sulfate ion) because they do not change during the reaction:

$$Mg(s) + Cu^{2+}(aq) \rightarrow Mg^{2+}(aq) + Cu(s)$$

Here, the magnesium atom is oxidised (loses electrons: $Mg \rightarrow Mg^{2+} + 2e^-$) and the copper ion is reduced (gains electrons: $Cu^{2+} + 2e^- \rightarrow Cu$). If a less reactive metal (like silver) is added to copper(II) sulfate, there is no change.

Displacement also happens with solid metal oxides, though this usually requires heating to overcome the activation energy. A famous example is the Thermite Reaction, where highly reactive aluminium powder displaces iron from iron(III) oxide:

$$Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s)$$

This reaction is extremely vigorous, producing intense heat, light, and molten iron.