In this tutorial, you’ll learn the basics of electrochemistry, including oxidation, reduction, galvanic cells, and applications of electrochemistry. We’ll also go over the fundamental electrochemistry equations and how to use them.
Electrochemistry is the field that studies the way chemical reactions are related to electricity. This includes chemical reactions that produce electricity like in batteries, chemical reactions that require electricity, and much more.
Topics Covered in Other articles
- Electrochemical Cells
- Understanding Redox Reactions
- How to Balance Redox Reactions
- What is the Equilibrium Constant?
- How to Write Net Ionic Equations
Oxidation: the process of an atom LOSING an electron
Reduction: the process of an atom GAINING an electron
Redox reaction: a chemical reaction where oxidation and reduction occur, meaning electrons are transferred between chemical species
Coulomb: the SI unit for electric charge (abbreviated C)
Voltage (aka Potential): the energy conveyed by electric charge (units of Joule/Coulomb)
Most chemical reactions are pretty simple from an electron’s perspective. Bonds may form or break, but generally, every atom will keep its electrons intact. This is not the case for redox reactions. “Redox” comes from the combination of “reduction” and “oxidation”. A redox reaction is one where electrons are lost by atoms or molecules and transferred to other ones. This transfer of electrons is the core of electrochemistry. For more information on redox reactions and how to balance them, see our linked tutorials.
Examples of Redox Reactions
2Na (s) + 2H2O (l) → 2NaOH (aq) + H2 (g)
Fe2+ (aq) + Ag+ (aq) → Fe3+ (aq) + Ag (s)
2Al (s) + 6H+ (aq) → 2Al3+ (aq) + 3H2 (g)
When a redox reaction is split into its reduction and oxidation half-reactions, electrons are written as products and reactants, respectively. They are called half-reactions because when “added” together, they make the overall redox reaction. Half reactions can often be considered independently.
A galvanic cell is a setup that facilitates redox reactions in a controlled and specific way that generates a current (flow of electrons). Voltaic cell, electrochemical cell, and battery are all different names for a galvanic cell. See our in-depth tutorial on galvanic cells here.
Voltaic Cell Diagram
Below is a diagram showing a typical voltaic cell.
Let’s break down the diagram and go over each part of the cell.
First, the “copper side” of the cell, called the reduction half-cell. It’s called this because it comprises half of the overall voltaic cell, and it is where electrons flow (meaning reduction occurs there). Incoming electrons reduce the copper ions in solution forming copper metal, which deposits on the surface of the copper electrode.
Next, the “zinc side” of the cell is called the oxidation half-cell. It’s called this because the zinc is giving up its electrons (oxidation). As electrons depart the zinc metal in the electrode, they flow over the wire to the reduction half-cell, and the left-over zinc ions enter the solution.
The wire connecting the two half-cells is what allows the electrons to flow. Without the wire, this setup does nothing, because there is nothing to allow the electrochemical reaction to progress. The wire generally includes a voltmeter, a device which measures the voltage of the cell.
Lastly, the salt bridge, the unsung hero of the galvanic cell. It contains spectator ion salts which balance out the charge buildup that forms as the reaction proceeds. Positively-charged zinc ions form in one cell and positively-charged copper ions leave the other as electrons flow. Unbalanced electrical charge is extremely unfavorable in chemistry, and as such the reaction cannot proceed without something to balance the charge. The salt bridge does just this, supplying positive ions to the reduction half-cell and negative ions to the oxidation half-cell. Without a salt bridge, galvanic cells do not work.
Spontaneity and equilibrium
You may recall from thermodynamics that some processes are spontaneous, or occur without the input of energy. ∆G (“delta-G”), the change in the Gibbs Free Energy of the system, represents spontaneity. When ∆G is negative, the energy of the system has gone down (favored). When ∆G is positive, the energy of the system has gone up (unfavored).
Electrochemists use ∆G less, because the voltage of a redox system represents the same thing – favorability. When the voltage of a reaction is POSITIVE, the reaction favors the forward direction and will occur spontaneously. When the voltage is NEGATIVE, the reaction disfavors the forward direction and will NOT occur spontaneously. This can be a bit confusing because the signs are different between voltage and ∆G. However, it can be easy to remember if you look at a reduction table. All the alkali metals and other easily-oxidized metals like zinc and iron are more favored to lose electrons than gain them. The reverse is true for nonmetals like fluorine and various metals like copper and silver that are hard to oxidize – these elements favor gaining electrons.
E0cell, the voltage, can be related mathematically to the free energy change ∆G and the equilibrium constant Keq. The equations below interrelate these three quantities:
R: the gas constant (in thermodynamic form), equal to 8.314 J/mol*K.
T: the temperature in Kelvin.
n: the number of electrons transferred in the reaction.
F: Faraday’s constant, equal to 96485 C/mol of electrons.
Potential of a cell
E0cell = E0red + E0ox = E0cathode – E0anode
For the cathode/anode equation above, use the reduction potentials for both. This means the potential at the anode may be negative. For the red/ox equation, use the reduction potential for the cathode and the opposite value for the anode (the oxidation potential).
The Nernst Equation
There is one more equation that is used frequently in electrochemistry, and that is the Nernst equation. Until this point, we’ve assumed that the concentrations of ions in solution doesn’t change. In other words, we’ve only been dealing with the initial potential of the cell in standard conditions. However, the reduced ions and/or anode in real-life batteries eventually run out. This is what happens when a battery dies – there are not enough cations remaining in the cathode solution, and/or too many cations in the anode solution for the reaction to proceed. The Nernst equation deals with non-equilibrium conditions like these. For an in-depth explanation on using the Nernst equation, see this article.
Electrolysis is the process of using an electric current to “force” a reaction to happen that otherwise would not. This technique is often used to decompose compounds. An example is the reaction 2H2O → 2H2 + O2 which generates hydrogen and oxygen gas, two very useful molecules, from water. This decomposition does not occur naturally, but readily proceeds with a relatively small voltage applied. Electrolysis sees wide industrial use to obtain various molecules from ores, seawater, and other common substances. For more information on electrolysis and the setup of electrolytic cells, see our tutorial.
In a normal galvanic cell, the cation in the reduction half-cell matches the cathode. That’s because those cations deposit on the surface as metal. The cation matches the cathode to keep things clean. However, the reaction can proceed just fine if the cation doesn’t match. The metal cation will convert to its neutral metal form and deposit on the surface of the cathode. This is called electroplating. This process protects metals under a layer of another, more resilient metal. For example, a layer of copper prevents easily-oxidized metals like iron from rusting. Electroplating is extremely common and a very useful technique for covering metal with another metal, which is very helpful in many applications.