Electrochemistry Introduction

Similar to Acid/Base reactions, it is useful to classify a large number of reactions as reduction – oxidation reactions. Redox reactions refer to reactions where one substance loses electrons (is oxidized) and another substance gains electrons (is reduced). The term oxidation comes for the reaction of many substances with oxygen. The substance loses electrons to oxygen (the substance is oxidized). Remember OILRIG (Oxidation Is Losing electrons, Reduction Is Graining electrons.) These reactions are also called electron transfer reactions. In a spontaneous redox reaction we can use the flow of electrons to do useful work such as run a flashlight or radio (remember when DG is negative it represents the maximum useful work that can be done by a reaction). Corrosion of metals are unwanted spontaneous reactions that can cost us money, when we repair corroded structures in time, and can cost lives when we don’t. We can use energy to make a non-spontaneous reaction run so as to make different substances. For example we can turn sodium chloride into the elements sodium and chlorine. Electrochemistry refers to the branch of chemistry that deals with the relationship between electricity and chemical reactions. We can use concepts from electrochemistry to produce useful power supplies such as batteries, to quickly measure quantities of substances, and to slow down spontaneous reactions we don’t want, such as the oxidation of the metals in our cars.

Remember oxidation numbers, where we assign a charge to atoms in elements, compounds, and ions. (please review pages 128-130 of Brown et al 9^th ed.). I wrote assign because in a compound like H₂O, we assign oxidation number of +1 to H and –2 to O, even though it is a molecular compound where electrons are shared, not transferred. Oxidation numbers are useful identifying redox reactions, because oxidation number change in redox reactions.

When we add hydrochloric acid to zinc metal, we get the following reaction

Molecular equation Zn (s) + 2 HCl (aq) = ZnCl₂ (aq) + H₂ (g)

Net Ionic equation Zn (s) + 2 H⁺ (aq) = Zn⁺² (aq) + H₂ (g)

Oxidation #s 0 1 2 0

Zinc loses 2 e^- and is oxidized

Each H⁺ gains an e^- and is reduced (its ox # is reduced from +1 to 0). The term oxidation comes from what most substances do when they react with O₂, they lose e^-.

The substance that is reduced, causes the other substance to be oxidized. Thus the substance reduced is called an oxidizing agent or oxidant. The substance oxidized caused another substance to be reduced and is called a reducing agent. It might help you to remember that oxygen, O₂, is a good oxidizing agent, (and is reduced in the process).

Balancing

All we are going to do is sum half reactions, so all electrons produced in an oxidation half reaction, are consumed in the reduction half reaction.

Zn (s) → Zn⁺² (aq) + 2 e^- Oxidation half reaction

+ 2 H⁺ (aq) + 2 e^- → H₂ (g) Reduction half reaction.

Zn (s) + 2 H⁺ (aq) → Zn⁺² (aq) + H₂ (g) The sum of two half reactions gives the net reaction.

Note that the charges need to balance in the half reactions, and the electrons provided by the oxidation half-reaction are consumed by the reduction half-reaction.

If you look at a piece of copper metal in a silver nitrate solution, you will see little branches of silver metal growing out of the copper wire. These little branches are called dendrites. You might also notice the colorless silver nitrate turning blue due to the oxidation of Cu (s) to Cu⁺². The two half reactions are;

Ag⁺(aq) + e^- → Ag (s) Reduction

Cu (s) → Cu⁺² (aq) + 2 e^- Oxidation

In order for this to balance, we need to multiple the top half-reaction by 2 so the electrons balance out.

2 Ag⁺ (aq) + 2 e^- → 2Ag (s) Reduction

+ Cu (s) → Cu⁺² (aq) + 2 e^- Oxidation

2 Ag⁺ + Cu → Cu⁺² + 2 Ag

This also works for reactions where water, an acid or base take place in the reaction. For instance, in analyzing the B group cations we converted Mn⁺² to MnO₂ using Ce⁺⁴. Mn in MnO₂ must have a +4 oxidation number, (why?) so Mn has been oxidized, and we are told that the Ce⁺⁴ is reduced to Ce⁺³. Appendix I in Moore et al and E in Brown et al 8th edition, has half reactions written as reduction reactions, so look for a half reaction where MnO₂ is reduced to Mn⁺² (the reverse of a reduction half reaction is an oxidation half reaction).

2Ce⁺⁴ + 2 e^- → 2 Ce⁺³ (check to make sure the charges balance on

+ Mn⁺² + 2 H₂O → MnO₂ + 4 H⁺ + 2 e^- both sides and electrons balance)

Mn⁺² + 2 Ce⁺⁴ + 2 H₂O → MnO₂ + 4 H⁺ + 2 Ce³

Voltaic (Galvanic) Cells

If I put a piece of zinc metal into a blue Cu²⁺ solution, the color gradually fades, and we see some black solid material produced on the zinc metal. The zinc metal is losing electrons to the copper cations as follows.

Zn (s) ® Zn²⁺ (aq) + 2 e^- oxidation (losing e^-)

Cu²⁺ +2 e^- ® Cu(s) reduction (gaining e^-)

Zn(s) + Cu²⁺(aq) ® Cu(s) + Zn²⁺ (aq)

If we separate these two half reactions, we can use the negative Gibbs free energy to perform electrical work. Here we set up an external pathway rather that directly between reactants.

Anode

Cathode

Salt Bridge

electron and current flow and ion flow across salt bridge