Electrode Potential: The Driving Force of Redox Reactions
The Basics of Electron Transfer
At its heart, chemistry is about electrons. Some atoms hold onto their electrons very tightly, while others are more willing to give them away. Imagine a playground where electrons are toys. Some children (atoms) are very possessive and will always try to take a toy (gain an electron). Other children are more generous and are happy to give a toy away (lose an electron). Electrode Potential is like a measure of how "greedy" or "generous" an atom or ion is with its electrons.
A high tendency to gain electrons means the species is easily reduced. We call this a high reduction potential. Conversely, a high tendency to lose electrons means the species is easily oxidized, which corresponds to a high oxidation potential. It is important to remember that oxidation potential is simply the negative of the reduction potential. If a species has a strong desire to be reduced, it forces its partner to be oxidized.
The Standard Hydrogen Electrode: The Ultimate Reference
To measure anything, you need a starting point, a zero on your ruler. For electrode potential, this zero is defined by the Standard Hydrogen Electrode (SHE)[3]. The SHE is based on the following half-reaction happening under very specific, or standard conditions[4]:
$2H^+_{(aq)} + 2e^- \rightarrow H_{2(g)}$
By international agreement, the electrode potential for this reaction is set to exactly 0.000 V (volts). Now, we can measure the potential of any other half-cell by connecting it to the SHE. If the other half-cell has a greater tendency to be reduced than the SHE, its potential is recorded as a positive number. If it has a lesser tendency, its potential is recorded as a negative number.
The Standard Electrode Potential ($E^{\circ}$)
When we measure the electrode potential of a half-reaction under standard conditions (1 M concentration for solutions, 1 atm pressure for gases, and a temperature of 298 K), we call it the Standard Electrode Potential, denoted as $E^{\circ}$. Scientists have compiled these values into a table, known as the Standard Electrode Potential Series or the Electrochemical Series. This table is incredibly powerful for predicting the direction of redox reactions.
A species with a more positive $E^{\circ}$ value has a stronger tendency to gain electrons and be reduced. A species with a more negative $E^{\circ}$ value has a stronger tendency to lose electrons and be oxidized.
| Half-Reaction (Reduction) | $E^{\circ}$ (V) |
|---|---|
| $F_{2(g)} + 2e^- \rightarrow 2F^-_{(aq)}$ | +2.87 |
| $Au^{3+}_{(aq)} + 3e^- \rightarrow Au_{(s)}$ | +1.50 |
| $Ag^+_{(aq)} + e^- \rightarrow Ag_{(s)}$ | +0.80 |
| $2H^+_{(aq)} + 2e^- \rightarrow H_{2(g)}$ (SHE) | 0.00 |
| $Zn^{2+}_{(aq)} + 2e^- \rightarrow Zn_{(s)}$ | -0.76 |
| $Na^+_{(aq)} + e^- \rightarrow Na_{(s)}$ | -2.71 |
| $Li^+_{(aq)} + e^- \rightarrow Li_{(s)}$ | -3.04 |
Predicting Spontaneity and Calculating Cell Voltage
The primary application of standard electrode potentials is predicting whether a redox reaction will occur spontaneously. The rule is simple: A redox reaction is spontaneous if the half-reaction with the more positive $E^{\circ}$ value undergoes reduction, and the half-reaction with the more negative $E^{\circ}$ value undergoes oxidation.
When you combine two half-cells to make a full electrochemical cell (like a battery), the overall cell potential, or voltage ($E^{\circ}_{cell}$), can be calculated. There are two common ways:
Method 1: $E^{\circ}_{cell} = E^{\circ}_{cathode} - E^{\circ}_{anode}$
Where the cathode is where reduction occurs (more positive E°) and the anode is where oxidation occurs (more negative E°).
Method 2: $E^{\circ}_{cell} = E^{\circ}_{reduction} + E^{\circ}_{oxidation}$
Since $E^{\circ}_{oxidation} = -E^{\circ}_{reduction}$ for the opposite half-reaction, this method is mathematically equivalent to the first.
A Concrete Example: The Lemon Battery
A classic school project that perfectly illustrates electrode potential is the lemon battery. You insert a strip of zinc (like a galvanized nail) and a copper coin into a lemon. When you connect the two metals with wires to a small light bulb or voltmeter, you create a simple electrochemical cell.
- The Zinc strip has a standard reduction potential of $Zn^{2+} + 2e^- \rightarrow Zn$ with $E^{\circ} = -0.76 V$. This negative value means zinc has a high tendency to lose electrons, so it acts as the anode and is oxidized: $Zn \rightarrow Zn^{2+} + 2e^-$.
- The Copper coin has a standard reduction potential of $Cu^{2+} + 2e^- \rightarrow Cu$ with $E^{\circ} = +0.34 V$. This positive value means copper ions have a tendency to gain electrons, so copper acts as the cathode where reduction occurs: $Cu^{2+} + 2e^- \rightarrow Cu$. The $Cu^{2+}$ ions come from the lemon's acidic juice.
The cell potential is calculated as:
$E^{\circ}_{cell} = E^{\circ}_{cathode} - E^{\circ}_{anode} = 0.34 V - (-0.76 V) = 1.10 V$
This positive voltage confirms the reaction is spontaneous. The electrons flow from the zinc anode (through the wire) to the copper cathode, creating an electric current that can, in theory, power a small device.
Important Questions
It provides a universal reference point of zero volts. Without a common reference, we could only say one metal is "more active" than another, but we couldn't assign a precise numerical value to its electrode potential. The SHE allows us to create a quantitative scale, the electrochemical series, which is essential for precise calculations in electrochemistry.
A negative $E^{\circ}_{cell}$ means the reaction is non-spontaneous under standard conditions. It will not happen on its own. However, it can be forced to occur by supplying electrical energy from an external source. This process is called electrolysis. For example, we use electrolysis to plate metals onto objects or to extract reactive metals like aluminum and sodium from their ores.
Corrosion, like the rusting of iron, is an electrochemical process. Metals with more negative (or less positive) reduction potentials are more likely to corrode (be oxidized). Iron ($E^{\circ} = -0.44 V$) corrodes easily in moist air. In contrast, gold ($E^{\circ} = +1.50 V$) does not corrode because it has a very high reduction potential, meaning it strongly resists being oxidized. This is why gold is found as a pure metal in nature and is considered a "noble" metal.
Footnote
[1] Oxidized/Reduced: Oxidation is the loss of electrons; Reduction is the gain of electrons (OIL RIG: Oxidation Is Loss, Reduction Is Gain).
[2] Redox: A portmanteau for Reduction-Oxidation reactions, which are chemical reactions involving the transfer of electrons between two species.
[3] Standard Hydrogen Electrode (SHE): A reference electrode which has a standard electrode potential defined to be exactly 0 V, against which all other electrode potentials are measured.
[4] Standard Conditions: For electrode potential measurements, this means a temperature of 298 K (25°C), a 1 M concentration for all aqueous solutions, and a 1 atm pressure for any gases involved.