Predicting a redox reaction

The electrochemical series

List the couples in order of $E^\circ$ and you get the electrochemical series. It is one of the most useful tables in chemistry.

Read it with two ideas:

• A couple with a very negative $E^\circ$ gives up electrons easily. Its reduced form, a metal such as potassium or magnesium, is a strong reducing agent.
• A couple with a very positive $E^\circ$ takes electrons easily. Its oxidised form, such as $\text{F}_2$, $\text{Cl}_2$ or $\text{Ag}^+$, is a strong oxidising agent.

So a reduced form can reduce the oxidised form of any couple with a higher $E^\circ$. In the picture, a metal can reduce the ions of every couple that stands above it.

Does a reaction happen?

Follow four steps:

1. Write the two half-reactions as reductions and look up their $E^\circ$.
2. The species that is reduced belongs to the couple with the higher $E^\circ$.
3. Work out $E^\circ = E^\circ(\text{reduced}) - E^\circ(\text{oxidised})$.
4. If it is positive, the reaction can happen under standard conditions.

Zinc in a copper solution. The copper ions are reduced and zinc is oxidised.

• $E^\circ = 0.34 - (-0.76) = +1.10\ \text{V}$. Positive, so zinc dissolves and copper deposits.

Copper in a zinc solution. Now the zinc ions would have to be reduced.

• $E^\circ = -0.76 - 0.34 = -1.10\ \text{V}$. Negative, so nothing happens.

Copper in a silver solution. The silver ions are reduced and copper is oxidised.

• $E^\circ = 0.80 - 0.34 = +0.46\ \text{V}$. Positive: $\text{Cu} + 2\text{Ag}^+ \rightarrow \text{Cu}^{2+} + 2\text{Ag}$. Silver crystals grow on a copper wire.

The pattern is simple. A metal displaces the ions of any metal that is higher in the series, and never the ones lower down.

Metals and acids

Hydrogen sits at $0.00\ \text{V}$. A metal can reduce $\text{H}^+$ to hydrogen gas if its $E^\circ$ is below zero.

• Magnesium, zinc and iron have negative $E^\circ$. They dissolve in dilute hydrochloric acid and release hydrogen.
• Copper and silver have positive $E^\circ$. Dilute hydrochloric acid does nothing to them.

Nitric acid is a different case. There the nitrate ion, not the hydrogen ion, is the oxidising agent. That is why copper does dissolve in nitric acid.

[Series figure above: the position of $\text{H}^+/\text{H}_2$ divides the metals that release hydrogen from those that do not.]

Cell potential and spontaneity

A positive $E^\circ_{cell}$ means a reaction can happen. The link to energy is:

Here $n$ is the number of moles of electrons in the balanced equation, and $F = 96500\ \text{C/mol}$. A positive $E^\circ$ gives a negative $\Delta G^\circ$.

For the Daniell cell, $n = 2$ and $E^\circ = 1.10\ \text{V}$:

• $\Delta G^\circ = -2 \times 96500 \times 1.10 = -212\,300\ \text{J/mol} = -212.3\ \text{kJ/mol}$

Take this link as given for now. In Class 12 you will see where it comes from, and how the potential changes when the concentrations are not $1\ \text{mol/L}$.

Remember one limit. $E^\circ$ says whether a reaction can happen. It does not say how fast. Some reactions with a positive $E^\circ$ are so slow that we never see them.

Redox in everyday life

Rusting. Iron rusts because it is oxidised by oxygen and water. A coat of zinc, called galvanising, protects it. Zinc has a lower $E^\circ$ ($-0.76\ \text{V}$) than iron ($-0.44\ \text{V}$), so zinc is oxidised first. The zinc is used up and the iron survives, even when the coat is scratched.

A coat of tin works differently. Tin has $E^\circ = -0.14\ \text{V}$, which is higher than iron's. Tin protects the iron only while the coat is whole. Once it is scratched, the iron is oxidised first and rusts faster.

Batteries. Any pair of couples with a positive cell potential can make a battery. A common dry cell uses a zinc case as its anode.

Practise with new numbers

Try it in Python

The program checks every strip against every solution, and works out a free energy change.

Common mistakes

• Thinking a more positive $E^\circ$ means a more reactive metal. It is the opposite: the more negative the $E^\circ$, the more readily a metal gives up electrons.
• Mixing up the order in the subtraction. The ions that are reduced come first.
• Leaving out the minus sign in $\Delta G^\circ = -nFE^\circ$.
• Using $E^\circ$ to say how fast a reaction goes. It says only whether the reaction can happen.

Remember these

• Lower $E^\circ$: stronger reducing agent. Higher $E^\circ$: stronger oxidising agent.
• A reaction can happen if $E^\circ = E^\circ(\text{reduced}) - E^\circ(\text{oxidised})$ is positive.
• Metals with $E^\circ < 0$ release hydrogen from dilute acid. Copper and silver do not.
• $\Delta G^\circ = -nFE^\circ$. A positive $E^\circ$ gives a negative $\Delta G^\circ$.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. Does iron dissolve in a solution of zinc sulphate?

2. Why does copper not release hydrogen from dilute hydrochloric acid?

3. A cell reaction has $E^\circ = +0.50\ \text{V}$ and moves 2 mol of electrons. Is $\Delta G^\circ$ positive or negative?

4. Why does a scratched zinc coat still protect iron?