IB Chemistry home > Syllabus 2016 > Redox processes > Relative reactivity

Syllabus ref: 9.1

'Reactivity' means the tendency for a substance to undergo a chemical reaction when placed in suitable circumstances. A substance that is chemically reactive in one situation may not necessarily be reactive in another. However, there are certainly trends in reactivity amongst similar compounds. This section illustrates reactivity in general terms by comparing the group I metals and the halogens (group 17).

Relative reactivity

The concept of reactivity means the relative activity of a chemical compound when reacting with other compounds. To provide a 'fair' test the compounds are usually compared with a specific 'standard' compound. For example the reactivity of metals may be assessed by reference to their behaviour with water.

The actual reactivity can be assessed in terms of the amount of energy released, or the relative speed, or violence, of the reaction. Taking the example above, the reaction of metals with water, it is observed that when potassium, sodium and lithium are placed separately in a bowl of cold water the observations serve to differentiate between the metals.

Lithium reacts rapidly, although the heat of the reaction is insufficient to melt the lithium. There are no sparks or flames and the lithium requires more time to dissolve. Hydrogen is evolved.

Sodium also melts into a ball and, although sparks are often seem, it doesn't ignite unless its motion is restricted in some way. It takes longer to finish reacting than potassium, but still dissolves rapidly. Hydrogen is evolved.

Potassium ignites immediately on contact with water. It melts into a ball and moves rapidly across the surface of the water. It reacts vigorously and dissolves in a very short time. Hydrogen is evolved and burns with the heat of the reaction.

We can say from the above observations that the order of reactivity is: potassium > sodium > lithium.

When metals react they do so by losing electrons. This process is called oxidation and the metals themselves are behaving as reducing agents. In the example above, the potassium is a stronger reducing agent than sodium and lithium, as it reduces the hydrogen ions in water faster and with a more exothermic reaction.

K(s) + H2O(l) KOH(aq) + H2(g)

The reaction can be broken down into two 'half' equations, one showing what happens to the species that gets reduced and the other to the species oxidised.

Potassium is oxidised to potassium ions:

K(s) K+(aq) + 1e

Hydrogen ions are reduced to hydrogen gas:

2H+(aq) + 2e H2(g)

Using reactions such as these we can arrive at a table of relative reducing and oxidising capacity for a group of compounds. Such a table or list is called a reactivity series.


Reactivity series of metals

It is possible to organise a group of similar chemicals that undergo either oxidation or reduction according to their relative reactivity. Oxidation (and reduction) is a competition for electrons. The oxidising species (agents) remove electrons from other species and can force them to become reducing agents (releasers of electrons)

A good example of this competition for electrons is the behaviour of metals. Metals always react by losing electrons (oxidation); they are reducing agents. However if a metal is in competition with metal ions of a different element, the more reactive metal can oblige the less reactive metal (in the form of ions) to accept electrons. This is called a displacement reaction.

Example: Zinc reacts with a solution containing copper ions. The zinc metal is more reactive than copper metal and so it can force the copper metal ions to accept electrons and become metal atoms.

Zn(s) Zn2+(aq) + 2e

Cu2+(aq) + 2e Cu(s)

The zinc metal passes its electrons to the copper ions. We observe that the zinc develops a pink layer of copper on its surface and the blue copper ion solution fades in colour.

We say that the zinc displaces the copper ions from solution.

Experimental observations

If we observe that there is a reaction between a metal and another metal ion in solution this tells us that the solid metal is more reactive than the metal of the dissolved metal ions.

  • Iron displaces copper from a solution of copper(II) sulfate
  • Copper displaces silver from a solution of silver nitrate

Given this information we can deduce that the most reactive of the three metals is iron, followed by copper, followed by silver. This allows us to arrange the metals into a reactivity series based on these specific reactions.

By convention, half-reactions for redox processes are written in a list as reductions:

Na+(aq) + 1e
Mg2+(aq) + 2e
Zn2+(aq) + 2e
Fe2+(aq) + 2e
2H+(aq) + 2e
Ag+(aq) + 1e

Ions of less reactive metals can 'steal' the electrons from more reactive metals. Hence in the table above magnesium metal can give its electrons to silver ions according to the reaction:

Ag+(aq) + Mg(s) Mg2+(aq) + Ag(s)

A species higher on the right hand side (a reducing agent) will react with an ion lower down on the left hand side (an oxidising agent).

Reduction of metal oxides by metals

metal A + metal B oxide metal A oxide + metal B

When a metal A is heated with a metal B oxide there will be a reaction if the free metal A is more reactive than the metal B of the metal B oxide. This is because the metal B in the metal B oxide is in the form of a metal ion - it has already lost electrons.

There is a competition between the metal ion (in the oxide) and the free metal for the electrons. The more reactive of the two metals will win the competition. Consequently if there is a reaction between a metal and a metal oxide then this tells us that the free metal is more reactive than the metal in the metal oxide.

Experimental observations

We can use this information to arrange the metals in order of reactivity

Reactivity of group 1 metals
most reactive
Copper least reactive

Sodium has the greatest electron releasing power (and conversely the copper ions - Cu2+ - would have the greatest electron attracting power)


Reactivity series of non-metals

Metals react by losing electrons - they are reducing agents. Non-metals react by gaining electrons - they are oxidising agents. In the same way that metals can be ordered in terms of reducing strength, the non-metals can be ordered in terms of their oxidising strength. The halogens are a typical example of a non-metal reactivity series.

Reactivity of the halogens
most reactive
Iodine least reactive

Fluorine is so reactive that we cannot isolate it in the laboratory very easily, as it reacts with both water and glass. As a result we don't usually deal with fluorine at pre-university level, but compare only the other three (astatine is very rare and radioactive)

Do not confuse this order of reactivity with that of the metals - these are non-metals, their reactivity is in terms of oxidising power - i.e. chlorine is the best oxidising agent out of chlorine, bromine and iodine.


Displacement reactions of halogens and halide ions

Chlorine displaces bromine from solutions containing bromide ions:

Cl2 + 2Br- Br2 + 2Cl-

In this reaction the chlorine is oxidising the bromide ions by removing an electron from them. Bromine is liberated from the solution and may be detected by its orange/red colour.

Bromine displaces iodine from solutions containing iodide ions:

Br2 + 2I- I2 + 2Br-

In this reaction the bromine is oxidising the iodide ions by removing an electron from them. Iodine is liberated from the solution and may be detected by its orange/brown colour which turns blue/black in the presence of starch indicator.

The results are explained by the order of oxidising power Cl2 > Br2 > I2. It is predictable, then, that chlorine will also displace iodine from a solution containing iodide ions:

Cl2 + 2I- I2 + 2Cl-