Electrochemical cells convert chemical energy to electrical energy

An electrochemical cell converts chemical energy into electrical energy through redox reactions. It is made up of two half-cells, with each containing an electrode submerged in a solution of ions.

These half-cells are linked by two key components:

1. An external circuit - This is a wire that allows the flow of electrons from one electrode to the other, creating an electric current. This current can be measured with a voltmeter.
2. A salt bridge - This bridge allows ions to transfer between the half-cells, maintaining charge balance. It's often a filter paper soaked in a neutral electrolyte like potassium nitrate.

The image below shows an electrochemical cell made by connecting a zinc half-cell with a copper half-cell.

In every electrochemical cell, two half-reactions occur: oxidation happens at the anode (the positive electrode), and reduction happens at the cathode (the negative electrode).

Types of half-cells

Electrochemical cells can be constructed from two main types of half-cells:

1. Metal/metal ion half-cell:

• Consists of a metal electrode in contact with a solution of its own ions.
• Example: A zinc electrode immersed in a solution of Zn²⁺ ions.
• The metal can either be oxidised to its ions or the ions can be reduced to the metal, depending on the metal's reactivity.

2. Ion-ion half-cell:

• Consists of two ions of the same element in different oxidation states, in contact with an inert platinum electrode.
• Example: Fe²⁺ and Fe³⁺ ions in solution with a platinum electrode.
• The platinum electrode provides an inert surface for the electron transfer between the ions of different oxidation states.

Zinc-copper electrochemical cell

A common example of an electrochemical cell is the zinc-copper cell, which is composed of a zinc half-cell and a copper half-cell:

1. At the zinc anode, zinc atoms are oxidised to Zn²⁺ ions, releasing electrons into the external circuit: Zn_(s) ➔ Zn²⁺_(aq) + 2e^-

2. At the copper cathode, Cu²⁺ ions are reduced to copper atoms by gaining electrons from the external circuit: Cu²⁺_(aq) + 2e^- ➔ Cu_(s)

Electrons move from the zinc, which is more reactive, to the copper, which is less reactive. The voltmeter measures the potential difference between the half-cells, known as the cell potential or electromotive force (e.m.f.), E_cell.

Electrode reactions are reversible

The reaction at each electrode in an electrochemical cell is reversible:

• At the zinc electrode: Zn_(s) ⇌ Zn²⁺_(aq) + 2e^-
• At the copper electrode: Cu²⁺_(aq) + 2e^- ⇌ Cu_(s)

The direction of each reaction depends on the ease with which the metal loses electrons (i.e., is oxidised).

This is quantified by the electrode potential of each half-cell:

• Metals that oxidise more readily have more negative electrode potentials.
• Metals that oxidise less readily have less negative or even positive electrode potentials.

In the zinc-copper cell, since zinc has a more negative electrode potential, it undergoes oxidation (the backward reaction). Copper has a more positive electrode potential, so it undergoes reduction (the forward reaction).

Shorthand notation for electrochemical cells

Electrochemical cells can be represented using a concise shorthand notation.

For the zinc-copper cell:

Zn_(s) | Zn²⁺_(aq) || Cu²⁺_(aq) | Cu_(s)

Key features of this notation include:

1. The half-cell with the more negative potential (i.e. Zn/Zn²⁺) is listed on the left.
2. A single vertical line indicates a phase boundary between the metal electrode and the aqueous ion.
3. A double vertical line denotes the salt bridge.
4. The oxidised species (i.e. Zn²⁺ and Cu²⁺) are written next to the salt bridge.

By following this convention, the cell potential (E^⦵_cell) is calculated from the standard electrode potentials (E^⦵) of the two half-cells:

$E_{\text{cell }}^{\ominus }=E_{\text{right-hand side }}^{\ominus }-E_{\text{left-hand side }}^{\ominus }$

For the zinc-copper cell:

E^⦵_cell = $+0.34-(-0.76)=+1.10\text{ V}$

Worked example 1 - Representing an iron-nickel cell

Iron metal reacts with nickel(II) ions to form iron(II) ions and nickel metal. Write the cell diagram for the electrochemical cell formed by this redox reaction.

Step 1: Write the full redox equation

Fe_(s) + Ni²⁺_(aq) ➔ Fe²⁺_(aq) + Ni_(s)

Step 2: Write the half-equations occuring at each electrode:

At cathode (reduction): Ni²⁺_(aq) + 2e^- ➔ Ni_(s)

At anode (oxidation): Fe_(s) ➔ Fe²⁺_(aq) + 2e^-

Step 3: Write the cell diagram:

Fe_(s) | Fe²⁺_(aq) || Ni²⁺_(aq) | Ni_(s)

Standard hydrogen electrode as reference

The standard hydrogen electrode acts as a universal reference for measuring electrode potentials. It features a platinum electrode immersed in a 1.00 mol dm^-3 solution of H⁺_(aq) ions at 298 K, with H₂ gas at a pressure of 100 kPa bubbling over the electrode.

The standard hydrogen electrode has a standard electrode potential (E^⦵) set at 0.00 V.

To determine the E^⦵ value of another half-cell, it's connected to the standard hydrogen electrode under standard conditions:

1. The temperature is kept at 298 K (25°C).
2. The pressure is kept at 100 kPa.
3. Solutions are at a concentration of 1.00 mol dm^-3 or are equimolar. The ^⦵ symbol beside E indicates these standard conditions are met.

In diagrams, the standard hydrogen electrode is always positioned on the left side, regardless of the relative E^⦵ values of the half-cells. The voltmeter's reading reflects the E^⦵ value of the other half-cell, which can be positive or negative, depending on the electron flow direction relative to the standard hydrogen electrode.

The standard electrode potential (E^⦵) of a half-cell is the voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode

$E_{\text{cell }}^{\ominus }=E_{\text{right-hand side }}^{\ominus }-E_{\text{left-hand side }}^{\ominus }$

$E_{\text{cell }}^{\ominus }=E_{\text{right-hand side }}^{\ominus }-0.00$

$E_{\text{cell }}^{\ominus }=E_{\text{right-hand side}}^{\ominus }$

Like any reversible reaction, the equilibrium position in a half-cell is influenced by changes in temperature, pressure, and concentration. Altering the equilibrium position leads to changes in the E_cell value.

To ensure consistency and comparability of electrode potential values across different cells, standard conditions are used when measuring electrode potentials.