9.3 - Reactivity
9.3.1: Deduce a reactivity series based upon the chemical behaviour of a group of oxidising and reducing agents. Examples include displacement reactions of metals and halogens. Standard electrode potentials or reduction potentials are not required.
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 then reducing agents. However if a metals is in competition with metal ions 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 coper on its surface and the blue copper ion solution fades in colour.
We say that the zinc displaces the copper ions from solution.
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 sulphate
- 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
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.
- Magnesium reacts with zinc oxide:- Mg + CuO MgO + Cu
- Sodium reacts with magnesium oxide: 2Na + MgO Na2O + Mg
- Zinc reacts with copper oxide:- Zn + CuO ZnO + Cu
We can use this information to arrange the metals in order of reactivity
Sodium has the greatest electron releasing power (and conversely the copper ions - Cu2+ - would have the greatest electron attracting power)
Predictions from reactivity series
Once a reactivity series is produced it can be used to predict reactions of pairs of reactant. For example in the table above it should be appreciated that magnesium will react with copper oxide reducing it to copper metal.
Any metal that is more reactive will react with compounds of less reactive metals.
Reactivity series involving 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
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.
1. Chlorine will displace 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
2. Bromine will displace 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.
It is predictable, then, that chlorine will also displace iodine from a solution containing iodide ions
9.3.2: Deduce the feasibility of a redox reaction from a given reactivity series. Students are not expected to recall a specific reactivity series.
Prediction of feasibility
Once a reactivity series is constructed depending on the reduction or oxidation ability of each species, we can use it to predict the feasibility (probability) of a reaction occurring between any two pairs of reactants.
If one of the substances is a reducing agent - i.e. it reacts by losing electrons then this must react with an oxidising agent - i.e. a species that gains electrons.
K K+ + 1e
|best reducing agents (left hand side species)|
Mg Mg2+ + 2e
Zn Zn2+ + 2e
Fe Fe2+ + 2e
Cu Cu2+ + 2e
H2 2H+ + 2e
2I- I2 + 2e
2Br- Br2 + 2e
2Cl- Cl2 + 2e
|best oxidising agents (right hand side species)|
Any species from the right hand side of one of the redox equilibria (the oxidising agent) can be predicted to react with any species above it on the left hand side of the redox equilibria (the reducing agent).
The species on the right hand side of the equilibria will gain electrons to go to the right hand side. They can only gain these electrons form species that are abopve them on the left hand side of the series.
We can therefore predict that chlorine (right hand side) will react with copper (left hand side) to form copper ions nad chloride ions according to the equation:
Cl2 + Cu Cu2+ + 2Cl-
Similarly, we can predict that iodide ions (left hand side) will NOT react with zinc ions (left hand side) as the zinc ions are poor oxidising agents and the iodide ions poor reducing agents.
Note: although a reaction may be predicted as feasible it does not mean that it will happen spontaneously. If the activation energy is high then it may need an extra "push" to get it going. - for example the reaction between chlorine and hydrogen needs a spark or ultraviolet light and then it is explosively fast.