9.4.1 The reactivity series and reactions of metals

The reactivity series

1. From the most reactive to the least reactive, the order is potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold.
2. Carbon and hydrogen are non-metals, but they are placed in the series as reference points for comparing metals.
3. The higher a metal is in the series, the more vigorously it reacts, so potassium reacts violently with water while gold hardly reacts with anything.
4. Metals above hydrogen react with dilute acids to give hydrogen gas, while metals below hydrogen do not.
5. The series is built from experimental evidence, by comparing how quickly metals react with cold water, steam and dilute acids.

Reactions with cold water

1. Potassium, sodium and calcium react with cold water to form a metal hydroxide and hydrogen.
2. Potassium
   1. Potassium floats, melts into a shiny ball, races across the surface and the hydrogen catches fire with a lilac flame.
   2. The reaction is so violent that the potassium may crackle and spit.
   3. The equation is $2K(\mathrm{s})+2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to 2KOH(\mathrm{aq})+{\mathrm{H}}_2(\mathrm{g})$.
3. Sodium
   1. Sodium floats, melts into a ball and moves quickly around the surface with fizzing until it disappears.
   2. The reaction is less violent than the reaction of potassium, and the hydrogen does not usually catch fire.
   3. The equation is $2Na(\mathrm{s})+2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to 2NaOH(\mathrm{aq})+{\mathrm{H}}_2(\mathrm{g})$.
4. Calcium
   1. Calcium sinks and fizzes steadily, giving off a stream of hydrogen bubbles.
   2. The water turns cloudy because calcium hydroxide is only slightly soluble, so some of it appears as a white solid.
   3. The equation is $\mathrm{Ca}(\mathrm{s})+2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to \mathrm{Ca}(\mathrm{OH}{)}_2(\mathrm{aq})+{\mathrm{H}}_2(\mathrm{g})$.
5. The solution left in each case is alkaline, because a metal hydroxide has formed.
6. The reactions become less vigorous from potassium to sodium to calcium, which matches their order in the series.
7. Magnesium reacts only extremely slowly with cold water, producing just a few bubbles over several days.

Magnesium and steam

1. Magnesium reacts very slowly with cold water but vigorously with steam.
2. Hot magnesium glows with a brilliant white light in steam and leaves a white powder.
3. The products are magnesium oxide and hydrogen, not magnesium hydroxide: $\mathrm{Mg}(\mathrm{s})+{\mathrm{H}}_2\mathrm{O}(\mathrm{g})\to \mathrm{MgO}(\mathrm{s})+{\mathrm{H}}_2(\mathrm{g})$.
4. Magnesium needs the higher temperature of steam to react quickly, which shows that it is less reactive than calcium.

Reactions with dilute hydrochloric acid

1. Metals above hydrogen react with dilute hydrochloric acid to form a metal chloride and hydrogen.
2. Magnesium reacts vigorously, with rapid effervescence, and the test tube becomes noticeably warm: $\mathrm{Mg}(\mathrm{s})+2HCl(\mathrm{aq})\to \mathrm{MgC}{\mathrm{l}}_2(\mathrm{aq})+{\mathrm{H}}_2(\mathrm{g})$.
3. Zinc reacts steadily, giving a moderate stream of bubbles: $\mathrm{Zn}(\mathrm{s})+2HCl(\mathrm{aq})\to \mathrm{ZnC}{\mathrm{l}}_2(\mathrm{aq})+{\mathrm{H}}_2(\mathrm{g})$.
4. Iron reacts slowly, giving only a few bubbles, and the solution gradually turns pale green as iron(II) chloride forms: $\mathrm{Fe}(\mathrm{s})+2HCl(\mathrm{aq})\to \mathrm{FeC}{\mathrm{l}}_2(\mathrm{aq})+{\mathrm{H}}_2(\mathrm{g})$.
5. Copper, silver and gold show no reaction with dilute hydrochloric acid.
6. These results are explained by the positions of the metals in the reactivity series.
   1. Magnesium, zinc and iron are above hydrogen, so they react with the acid and release hydrogen.
   2. The further above hydrogen a metal is, the more vigorous its reaction, which is why magnesium reacts faster than zinc and zinc reacts faster than iron.
   3. Copper, silver and gold are below hydrogen, so they cannot release hydrogen from the acid and no reaction happens.

Deducing an order of reactivity

1. You may be given experimental results for metals you have not met before and asked to place them in order.
2. Compare results only when the conditions are the same, such as the same reagent, concentration, temperature and size of metal pieces.
3. A metal that reacts is more reactive than a metal that does not react under the same conditions.
4. When two metals both react, the one that gives faster bubbling or a larger temperature rise is more reactive.
5. A metal that reacts with cold water is more reactive than one that reacts only with steam.
6. A metal that does not react with dilute acid must be below hydrogen.
7. Write each result as a pair, such as "W is more reactive than X", and then join the pairs into a single list.

9.4.2 Reactivity, displacement and the oxide layer of aluminium

Reactivity and positive ions

1. When a metal reacts, its atoms lose electrons from their outer shell and become positive ions.
   1. Magnesium atoms each lose two electrons: $\mathrm{Mg}\to \mathrm{M}{\mathrm{g}}^{2+}+2{\mathrm{e}}^{-}$.
   2. Copper atoms also lose two electrons: $\mathrm{Cu}\to \mathrm{C}{\mathrm{u}}^{2+}+2{\mathrm{e}}^{-}$.
2. A more reactive metal has a greater tendency to lose electrons and form positive ions.
3. Magnesium loses its electrons far more readily than copper, so magnesium is much higher in the reactivity series.
4. The ions of a less reactive metal have a greater tendency to gain electrons and turn back into metal atoms.
5. This difference between two metals is what drives a displacement reaction.

Displacement reactions

1. In a metal displacement reaction, a more reactive metal is placed in a solution that contains the aqueous ions of a less reactive metal.
2. Atoms of the more reactive metal lose electrons and pass into solution as positive ions.
3. Ions of the less reactive metal gain these electrons and come out of solution as a coating of solid metal.
4. Zinc displaces copper from copper(II) sulfate solution: $\mathrm{Zn}(\mathrm{s})+\mathrm{CuS}{\mathrm{O}}_4(\mathrm{aq})\to \mathrm{ZnS}{\mathrm{O}}_4(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})$.
5. The sulfate ions are unchanged, so they are spectator ions and are left out of the ionic equation: $\mathrm{Zn}(\mathrm{s})+\mathrm{C}{\mathrm{u}}^{2+}(\mathrm{aq})\to \mathrm{Z}{\mathrm{n}}^{2+}(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})$.
6. During this reaction the blue colour of the solution fades, a red-brown coating of copper forms on the zinc, and the temperature rises.
7. Zinc atoms lose electrons, so zinc is oxidised, while copper(II) ions gain electrons, so they are reduced.
8. A displacement reaction is therefore a redox reaction.
9. If the less reactive metal is placed in a solution of the more reactive metal's ions, no reaction happens, so copper in zinc sulfate solution shows no change.
10. The further apart two metals are in the reactivity series, the larger the temperature rise in their displacement reaction usually is.

Predicting displacement for five metals

1. The order of magnesium, zinc, iron, copper and silver, from most to least reactive, is $\mathrm{Mg}>\mathrm{Zn}>\mathrm{Fe}>\mathrm{Cu}>\mathrm{Ag}$.
2. Each metal displaces the ions of every metal below it and none of the ions of the metals above it.
   1. Magnesium displaces zinc, iron(II), copper(II) and silver ions.
   2. Zinc displaces iron(II), copper(II) and silver ions.
   3. Iron displaces copper(II) and silver ions.
   4. Copper displaces silver ions only.
   5. Silver displaces none of these ions.
3. Iron in copper(II) sulfate solution turns the blue solution pale green and becomes coated with brown copper: $\mathrm{Fe}(\mathrm{s})+\mathrm{C}{\mathrm{u}}^{2+}(\mathrm{aq})\to \mathrm{F}{\mathrm{e}}^{2+}(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})$.
4. Copper in silver nitrate solution grows shiny grey crystals of silver, while the colourless solution slowly turns blue: $\mathrm{Cu}(\mathrm{s})+2A{\mathrm{g}}^{+}(\mathrm{aq})\to \mathrm{C}{\mathrm{u}}^{2+}(\mathrm{aq})+2Ag(\mathrm{s})$.
5. Magnesium in iron(II) sulfate solution becomes coated with dark grey iron as the pale green colour fades: $\mathrm{Mg}(\mathrm{s})+\mathrm{F}{\mathrm{e}}^{2+}(\mathrm{aq})\to \mathrm{M}{\mathrm{g}}^{2+}(\mathrm{aq})+\mathrm{Fe}(\mathrm{s})$.

The oxide layer on aluminium

1. Aluminium is high in the reactivity series, just below magnesium, yet in everyday use it often seems unreactive.
2. An aluminium saucepan does not fizz when water is boiled in it, and aluminium foil placed in copper(II) sulfate solution shows little or no change.
3. This is because aluminium is always covered by a very thin, tough layer of aluminium oxide, $\mathrm{A}{\mathrm{l}}_2{\mathrm{O}}_3$.
4. The layer forms as soon as fresh aluminium meets oxygen in the air: $4Al(\mathrm{s})+3{\mathrm{O}}_2(\mathrm{g})\to 2A{\mathrm{l}}_2{\mathrm{O}}_3(\mathrm{s})$.
5. The oxide layer is unreactive, it sticks firmly to the metal and it has no gaps, so water, acids and aqueous ions cannot reach the aluminium underneath.
6. If the surface is scratched, a new layer of oxide forms almost at once, so the protection repairs itself.
7. When the oxide layer is removed, the aluminium underneath reacts as vigorously as its position in the series predicts.
   1. In dilute hydrochloric acid, aluminium reacts slowly at first and then fizzes vigorously once the acid has dissolved the oxide layer.
8. The apparent unreactivity of aluminium is caused by its oxide layer, not by a low tendency to form positive ions.
9. The same oxide layer explains why aluminium resists corrosion and can be used for food containers.