Predicting reaction tendencies
Watch the following two videos.
The first shows an iron nail being placed in a solution of copper(II) sulfate.
The second shows a gold ring being placed into the same solution.
Text alternative to iron nail and copper sulfate video
The iron nail becomes coated in copper.
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Text alternative to gold ring and copper sulfate video
There is no reaction between the gold and the copper(II) sulfate.
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So why is there a reaction only in the first video?
Could we have predicted this behaviour using the reduction potentials?
The answer to the second question is yes, and the following section explains how this can be done.
Redox reactions and electrons
It seems a long time since we mentioned electrons, but understanding what is happening to electrons is vital to understanding redox reactions. So let's start by considering the reaction shown above in terms of electrons.
The two half-reactions involved in this process are shown below. The species present at the start of the reaction are Fe and Cu2+ ions.
(Note that these are listed in the order shown on the reduction potential list. It is good to do this because it helps with the explanation described later.)
Cu2+ + 2 e- 
Cu
Fe2+ + 2 e- 
Fe
The reaction that you saw in the video was copper metal being formed. This means the first reaction is moving in the forward direction, ie it is a reduction reaction. Therefore the iron reaction must move in the opposite direction.
The question to ask is: is iron metal (Fe) a strong enough reducing agent to make this happen?
Now look at your list of reduction potentials.
Is Fe a stronger reducing agent than Cu?
Yes. Fe is lower down the list (lower reduction potential) than Cu so it is stronger reducing agent than Cu. This means that Fe will give up its electrons to the Cu2+ ions, and the reaction will proceed spontaneously .
The overall reaction occurring is shown below.
Cu2+(aq) + Fe(s) 
Cu(s) + Fe2+(aq)
The two half-reactions involved in this process are shown below. The species present at the start of the reaction are Au and Cu2+ ions.
(Note that again these are listed in the order shown on the reduction potential list.)
Au3+ + 3 e- 
Au
Cu2+ + 2 e- 
Cu
If a reaction is to occur, the copper must be reduced. Is gold (Au) a strong enough reducing agent to make this happen?
No. Au is higher up the list than Cu so it is a weaker reducing agent than Cu. This means that Au will not be powerful enough to give up its electrons to the Cu2+ ions, and the reaction will not occur.
Now watch the videos below that work through two more examples of predicting reaction tendency.
Text alternative for the copper and hydrochloric acid video
Will copper metal react with dilute hydrochloric acid?
The reaction of a dilute acid with a metal can be described using the following half-equation.
2 H+(aq) + 2 e- 
H2(g)
That is, the hydrogen ion is being reduced and is acting as an oxidising agent.
When copper metal reacts it will be oxidised to produce copper(II) ions, so the relevant half-equation is shown below.
Cu2+(aq) + 2 e- 
Cu(s)
However, by looking at the table of standard reduction potentials, Cu2+ is a stronger oxidising agent than H+. Therefore, this reaction will not occur spontaneously. This is because hydrogen ions are not a strong enough oxidising agent to oxidise the copper metal to copper two plus ions.
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Text alternative for the chlorine and silver video
Will chlorine react with silver metal?
When chlorine reacts it will be reduced to form chloride ions
so the following is the relevant half-equation.
Cl2(g) + 2 e- 
2 Cl-(aq)
When silver metal reacts it will be oxidised to produce silver ions, so the following is the relevant half-equation.
Ag+(aq) + e- 
Ag(s)
That is, the chlorine ion is being reduced and is acting as an oxidising agent.
By looking at the table of standard reduction potentials, Cl2 is a stronger oxidising agent than Ag+. Therefore, this reaction will occur spontaneously. The chlorine half-equation proceeds in a forward reaction and the silver half-equation proceeds in the reverse direction.
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Predicting reaction tendency example
Below is a typical question that involves predicting reaction tendency.
Will potassium dichromate solution oxidise iodide ions?
The problem is solved in the following interaction.
Begin by selecting the two relevant half-equations by clicking on the left hand boxes. Then follow the prompts.
Text alternative to the predicting reaction tendency interaction
The interaction shows the process by which potassium dichromate solution oxidises iodide ions. It shows, by using information from the data sheet that the dichromate ions are a stronger oxidising agent than iodine, so the reaction will occur.
The two half-equations involved are
Reduction -
Cr2O72-+ 14 H++ 6 e-
2 Cr3++ 7 H2O
Oxidation -
I2 + 2 e-
2 I-
and the overall reaction will be:
Cr2O72- + 14 H+ + 3 I2 → 2 Cr3+ + 7 H2O + 2 I-
Click here to open the interaction in a new window
There are different ways to predict whether redox reactions will occur. The method shown above involves comparing strengths of oxidising agents and reducing agents involved in the reaction being investigated. By using this method you are consolidating your understanding of redox reactions while predicting whether reactions will take place.
There are some other 'shortcut' methods of predicting reaction tendency which are explained briefly below.
For a reaction to occur spontaneously, the oxidant in the reaction must have a more positive reduction potential than the reactant.
Example: Iron and copper(II) sulfate reaction
Cu2+ + 2 e- 
Cu (E° = +0.34 V)
Fe2+ + 2 e- 
Fe (E° = -0.44 V)
The reaction will occur because the reduction potential for the copper is higher than the reduction potential for the iron reaction.
Because the list of reduction potentials is ordered from highest down to lowest, it is possible to predict reaction tendency without looking at the actual values of the reduction potentials.
For reaction to occur spontaneously, the reduction (forward) reaction must be above the oxidation (reverse) on the list. Because the top reaction moves forward and the bottom reaction moves backwards, this can be called the clockwise rule because the reactions seem to move in a clockwise direction.
Example 1: Iron and copper(II) sulfate reaction
Click on the image to enlarge
The reaction satisfies the clockwise rule and the reaction should occur.
Example 2: Gold and copper(II) sulfate
Click on the image to enlarge
The equations seem to be trying to move in an 'anticlockwise' direction and so the reaction will not happen.
Although these shortcut methods are quick to use, they do not explain the reasons behind whether a reaction will occur. To do this you should go back to compare the strengths of the reactants as oxidising and reducing agents.
Limitations of the standard reduction potentials
- Remember that these are predictions based on chemicals under standard conditions. Even reactions that are predicted to occur spontaneously may not react. An example of this is where the concentrations are not high enough to allow a reaction to occur. (Remember the situation with the hydrogen peroxide bleach in the first scenario in this topic.)
- If there is only a small difference in reduction potential, the difference in oxidising or reducing strength may also not be enough to cause a reaction to happen.
Text Reference: Read Section 9.5 and answer Review Exercise 9.5
Consolidate your predicting reaction tendency skills by completing the worksheet 
.
