Establishing and Explaining the Reactivity Series

What metal reactivity means

Magnesium bubbles rapidly in dilute hydrochloric acid, whereas copper shows no reaction. These observations reveal a difference in reactivity: how readily the metals undergo chemical reactions.

When a metal reacts, its atoms lose electrons and form positive ions. Electrons have a negative charge, so losing them leaves the atom positively charged. The more readily a metal loses electrons to form its positive ion, the more reactive it is.

For example, a magnesium atom loses two electrons:

$$\mathrm{Mg \rightarrow Mg^{2+} + 2e^-}$$

This loss of electrons is called oxidation. Potassium and sodium form $\mathrm{K^+}$ and $\mathrm{Na^+}$ ions; magnesium forms $\mathrm{M}{\mathrm{g}}^{2+}$; aluminium forms . The size of the ion's charge does not tell you the metal's reactivity: potassium is more reactive than aluminium even though each potassium atom loses only one electron.

Establishing an order from reactions with water

A reactivity series places elements in order from most to least reactive. We can build it from experimental observations rather than simply memorising a list.

Metals that react with cold water produce a metal hydroxide and hydrogen:

metal + water → metal hydroxide + hydrogen

Potassium reacts very rapidly, sodium reacts rapidly, and calcium reacts less vigorously. Magnesium reacts very slowly with cold water, with bubbles gradually forming on its surface. Aluminium, zinc, iron, lead and copper usually show no visible reaction with cold water under ordinary laboratory conditions.

These observations support the order potassium > sodium > calcium > magnesium, where > means ‘more reactive than’. However, water alone cannot put all the metals showing no visible reaction into order.

The bubbles are hydrogen gas. Hydroxide formation makes the water alkaline, so universal indicator changes from green towards blue or purple. For calcium, the balanced equation is:

$$\mathrm{Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2}$$

Cold water and steam are different reaction conditions. For example, heated magnesium reacts with steam to form magnesium oxide, rather than magnesium hydroxide:

$$\mathrm{Mg + H_2O(g) \rightarrow MgO + H_2}$$

Using dilute acids to distinguish more metals

Metals above hydrogen in the reactivity series can react with dilute hydrochloric acid to produce a salt and hydrogen:

metal + dilute acid → salt + hydrogen

With hydrochloric acid, the salt is a chloride. Magnesium, for example, produces magnesium chloride. In comparable conditions, magnesium produces hydrogen faster than zinc, and zinc faster than iron. This supports magnesium > zinc > iron.

Hydrogen is included in the series as a reference point, although it is not a metal. Copper and silver are below hydrogen and do not release hydrogen from dilute hydrochloric acid.

Combining evidence from different reactions gives this order for the required elements:

Aluminium illustrates why observations need interpretation. Its oxide layer separates the underlying metal from the reactant. Little bubbling does not mean that aluminium belongs below iron in the series.

Comparing reactions fairly and safely

For a small-scale acid comparison, use test tubes in a rack, measured equal volumes of dilute hydrochloric acid, and small pieces of suitable metals such as magnesium, zinc and iron. Keep acid concentration and starting temperature the same, and use comparable exposed surface areas. A powder can react faster than a strip simply because more metal is exposed.

Observe how quickly bubbles form over the same time interval. A gas syringe can give a clearer comparison by measuring the volume of hydrogen produced in a fixed time. Faster hydrogen production suggests greater reactivity when the conditions are comparable; total gas produced also depends on the quantities of reactants.

Wear eye protection and use small quantities under teacher supervision. Potassium and sodium reacting with water are teacher demonstrations, not pupil test-tube experiments. Do not test them with dilute acid: their reactions are dangerously vigorous. Keep flames away from hydrogen-producing reactions except during a controlled gas test. A small collected sample of hydrogen gives a squeaky pop with a lighted splint.

When bubbling stops, gas production has ceased. This may mean that the metal or acid has been used up, rather than that both have disappeared. Handle any remaining metal and solution as instructed; collect potentially harmful metal-containing waste rather than pouring it away indiscriminately.

Displacement reactions provide another test

A more reactive metal displaces a less reactive metal from its compound. This is particularly useful for comparing metals that cannot be distinguished using cold water.

Place zinc in blue copper sulfate solution. A copper deposit forms on the zinc and the blue colour fades as copper ions are removed from solution:

zinc + copper sulfate → zinc sulfate + copper

$$\mathrm{Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)}$$

The reaction establishes that zinc is more reactive than copper. Copper placed in zinc sulfate solution does not displace zinc.

To interpret a results table, identify the solid metal added and the metal present as ions in the solution. A displacement means the added metal is more reactive. Link comparisons together: zinc displaces iron, and iron displaces copper, so zinc > iron > copper. Do not force a complete ranking if the available observations cannot distinguish two metals.

At particle level, zinc atoms lose electrons to become $\mathrm{Zn^{2+}}$ ions. Copper ions gain those electrons to become copper atoms. Electron gain is reduction, so oxidation and reduction occur together.

Writing formulae and balanced equations

Chemical symbols distinguish elements: calcium is Ca, copper is Cu and lead is Pb. A compound's formula shows the ratio of its atoms or ions. In an ionic compound, the positive and negative charges must balance overall.

Magnesium chloride contains $\mathrm{Mg^{2+}}$ and $\mathrm{Cl^-}$. Two chloride ions balance one magnesium ion, giving $\mathrm{MgCl_2}$. Aluminium chloride is . Calcium hydroxide is because one needs two ions; the brackets show that the whole OH group occurs twice.

Atoms are rearranged during reactions, not created or destroyed. This conservation of mass means a balanced equation has the same number of atoms of each element on both sides.

Start the magnesium–hydrochloric acid reaction with the correct formulae:

$$\mathrm{Mg + HCl \rightarrow MgCl_2 + H_2}$$

There are two chlorine atoms and two hydrogen atoms on the right, but only one of each on the left. Put 2 in front of HCl:

$$\mathrm{Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)}$$

Now both sides contain one magnesium, two chlorine and two hydrogen atoms. State symbols show solid (s), liquid (l), gas (g) and dissolved in water (aq).

For aluminium, balancing requires larger coefficients:

$$\mathrm{2Al + 6HCl \rightarrow 2AlCl_3 + 3H_2}$$

Both sides contain two aluminium, six chlorine and six hydrogen atoms. Coefficients multiply the entire formula that follows them.

Higher tier: writing ionic equations

An ionic equation leaves out ions that do not change during the reaction. These unchanged particles are spectator ions.

For magnesium reacting with hydrochloric acid, begin with the balanced symbol equation above. Write the dissolved acid and salt as their ions:

$$\mathrm{Mg(s) + 2H^+(aq) + 2Cl^-(aq) \rightarrow Mg^{2+}(aq) + 2Cl^-(aq) + H_2(g)}$$

The chloride ions appear unchanged on both sides. Cancel them to obtain:

$$\mathrm{Mg(s) + 2H^+(aq) \rightarrow Mg^{2+}(aq) + H_2(g)}$$

Leave the solid metal and hydrogen molecules intact; they are not dissolved ions. Check both atoms and charge: there is one magnesium and two hydrogen atoms on each side, and the total charge is +2 on each side.

The same method gives the ionic equation for zinc displacing copper:

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

Sulfate ions are spectators. This equation makes the chemical change clear: zinc becomes positive ions while copper ions become solid copper.

Reactivity and order

More reactive metals lose electrons more readily to form positive ions.

K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Ag

Hydrogen is a non-metal reference point. Aluminium's oxide layer can hide its reactivity.

Reading experimental evidence

• Cold water: reacting metals form hydroxide + hydrogen. K reacts faster than Na, then Ca; Mg is very slow.
• Dilute hydrochloric acid: metals above H can form chloride + hydrogen. Mg reacts faster than Zn, then Fe. Cu and Ag do not produce hydrogen.
• Displacement: the added metal must be more reactive than the metal in the compound.
• Compare rates under matched conditions. No visible change may not establish a metal's true position.

Formulae and balancing

Balance ionic charges to write formulae: $\mathrm{MgCl_2}$, , .