1. Metal + oxygen  metal oxide
Reactions of metals
Reactions with oxygen (combustion)
All metals form oxides except Ag, Au and Pt
Metal + oxygen  metal oxide
e.g. 2Mg + O2  2MgO
Tendency to form metal oxides:
Li, Na, K, Ca, Ba (react at room temp)
Mg, Al, Fe, Zn (react slowly at room temp, vigorously when heated)
Sn, Pb, Cu (react slowly and only when heated)
heat

2. Reactions of metals Reactions with water
Reactive metals react with water or steam
Metal + water  metal hydroxide + hydrogen gas
e.g. Na + 2H2O  2NaOH + H2
Metal + steam  metal oxide + hydrogen gas
e.g. Zn + H2O  ZnO + H2
Relative reactivity:
Li, Na, K, Ca, Ba (react with water at room temp)
Mg, Al, Zn, Fe (react with steam at high temp)
Sn, Pb, Cu, Ag, Au, Pt (do not react)

3. Metal + acid  salt + hydrogen gas
Reactions of metals
Reactions with dilute acid
More metals react with acid than water
Metal + acid  salt + hydrogen gas
Zn + 2HCl  ZnCl2 + H2
Relative reactivity:
Li, Na, K, Ca, Mg, Al, Zn, Fe, Co, Ni (react readily)
Sn, Pb (slow to react without heat)
Ag, Hg, Pt, Au (do not react)

4. Reactions of metals
Based on the ease of reactions with oxygen, water and acids, metals can be organised in order of reactivity, known as an activity series.
Activity series for metals:
K>Na>Ba>Ca>Mg>Al>Zn>Fe>Sn>Pb>Cu>Ag>Hg>Pt>Au
most reactive least reactive
Grp 1>Grp 2> Grp 3>some TM (Zn, Fe)>Grp 4>more TM
N.B. TM = transition metals

5. Oxidation-Reduction (REDOX)
The reactions of metals with oxygen, water and acids involve the metals losing electrons to form +ve metal ions.
When an atom loses one or more electrons, it is oxidised. If an atom gains electrons, it is reduced. Therefore:
Oxidation is loss of e-
Reduction is gain of e-
In any equation, there is no overall loss or gain of e-. Therefore, oxidation and reduction occur simultaneously and are known as redox reactions.
Originally, oxidation was gain of oxygen or loss of hydrogen and reduction was gain of hydrogen or loss of oxygen. However, this does not apply to redox reactions that do not include these elements. Therefore, loss/gain of electrons was adopted.

6. Rust as a redox reaction
In all metal corrosion reactions, the metal is oxidised to form a positive metal ion (i.e. loses electrons).
The more reactive the metal, the more likely the metal is to be oxidised.
Iron is oxidised by oxygen in the presence of water to form rust. The overall reaction is:
4Fe(s) + 3O2(g) + 2H2O(l) → 2Fe2O3 . xH2O(s) (rusting)
Note: x is a value from 1-3 indicating waters of hydration

7. Rust as a redox reaction
The two initial reactions involved in (wet corrosion)rusting are: Fe(s) → Fe 2+(aq) + 2e– (oxidation) and O2(g) + 2H2O(l) + 4e– → 4OH– (aq) (reduction) Iron(II) reacts with hydroxide to form the green precipitate, iron(II) hydroxide Fe 2+(aq) + 2OH–(aq) → Fe(OH)2(s) (green rust)

8. Rust as a redox reaction
Further exposure to moisture and oxygen leads to the oxidation of iron(II)hydroxide to red-brown iron(III)hydroxide 4Fe(OH)2 (s) + 2H2O(l) + O2(g) → 4Fe(OH)3(s) Iron(III)hydroxide then dehydrates to form rust 2Fe(OH)3 (s) → Fe2O3.xH2O (s) (rust)

9. Corrosion
Some metals, like aluminium form a protective oxide layer. Iron, however, produces hydrated iron oxide in the presence of water and air. This product flakes off and exposes more of the iron to further oxidation.

10. Rusting is a destructive process