 Question  Question: If you have a variety of metals and you want to determine which is the most reactive and which is the least reactive, how would you do it?  Answer  Answer: Try to react them with other substances and the ones that have the most vigourous reactions are the most reactive

1. Reactions with oxygen (combustion) All metals form oxides except Ag, Au and Pt Metal + oxygen  metal oxide e.g. 2Mg + O 2  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 with water Reactive metals react with water or steam Metal + water  metal hydroxide + hydrogen gas e.g. Na + 2H 2 O  2NaOH + H 2 Metal + steam  metal oxide + hydrogen gas e.g. Zn + H 2 O  ZnO + H 2 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. Reactions with dilute acid More metals react with acid than water Metal + acid  salt + hydrogen gas Zn + 2HCl  ZnCl 2 + H 2 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)

activity series Based on the ease of reactions with 1.oxygen, 2.water and 3.acids, metals can be organised in order of reactivity, known as an activity series. Activity series for metals 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

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.

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) + 3O 2 (g) + 2H 2 O(l) → 2Fe 2 O 3. xH 2 O(s) (rusting) Note: x is a value from 1-3 indicating waters of hydration

The two initial reactions involved in (wet corrosion)rusting are: Fe (s) → Fe 2+ (aq) + 2e– (oxidation) and O 2(g) + 2H 2 O (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)

Further exposure to moisture and oxygen leads to the oxidation of iron(II)hydroxide to red-brown iron(III)hydroxide 4Fe(OH) 2 (s) + 2H 2 O (l) + O 2(g) → 4Fe(OH) 3(s) Iron(III)hydroxide then dehydrates to form rust 2Fe(OH) 3 (s) → Fe 2 O 3.xH 2 O (s) (rust)

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.