Chemistry (9620) | Electrode potentials and electrochemical cells

Electrode Potentials and Electrochemical Cells: The Chemistry of Batteries!

Welcome to one of the most practical and exciting areas of chemistry: Electrochemistry! Don't worry if this topic feels daunting; essentially, we are just studying redox reactions that happen indirectly, forcing electrons to flow through a wire.
This is the science behind every battery, smartphone, and electric car in the world! Understanding this chapter helps us harness chemical energy to do useful electrical work.

Quick Review: Redox Reactions

Remember AS Chemistry? Redox means Reduction and Oxidation.

• Oxidation is Loss of Electrons (OIL)
• Reduction is Gain of Electrons (RIG)

In an electrochemical cell, we physically separate the oxidation process from the reduction process. This forces the electrons to travel through an external circuit, creating an electric current.

3.1.9.1 Electrode Potentials and Cells

What is a Half-Cell?

A simple redox reaction can be split into two parts: an oxidation half-reaction and a reduction half-reaction. When we physically separate these two parts, we create a half-cell.

A half-cell consists of a metal (the electrode) in contact with a solution containing its own ions. For example, a zinc half-cell consists of a zinc rod immersed in a \(1.00 \text{ mol dm}^{-3}\) solution of \(\text{Zn}^{2+}(\text{aq})\).

At the surface of the electrode, an equilibrium is set up:

\(\text{M}^{n+}(\text{aq}) + n\text{e}^- \rightleftharpoons \text{M}(\text{s})\)

The position of this equilibrium determines how easily the metal ions gain electrons (reduction), and thus determines the Electrode Potential (\(E\)).

IUPAC Convention for Half-Equations

The International Union of Pure and Applied Chemistry (IUPAC) sets a strict rule for writing half-equations for electrodes: they must always be written as reduction reactions (electron gain).
This is because the standard potential, \(E^\ominus\), is a measure of the tendency for reduction to occur.

Example:

• Copper half-equation: \(\text{Cu}^{2+}(\text{aq}) + 2\text{e}^- \rightarrow \text{Cu}(\text{s})\)
• Zinc half-equation: \(\text{Zn}^{2+}(\text{aq}) + 2\text{e}^- \rightarrow \text{Zn}(\text{s})\)

(Note: Even if zinc is oxidised in the overall cell, its half-equation is written as a reduction for comparison purposes).

Standard Electrode Potential (\(E^\ominus\))

The voltage measured for a half-cell depends heavily on concentration and temperature. To compare different half-cells fairly, we must measure them under Standard Conditions. When measured under these specific conditions, the potential is called the Standard Electrode Potential (\(E^\ominus\)).

Standard Conditions (Must be Known):

• Temperature: \(298 \text{ K}\) (\(25^\circ \text{C}\))
• Pressure: \(100 \text{ kPa}\) (for any gases involved)
• Concentration: \(1.00 \text{ mol dm}^{-3}\) (for all ions involved)

The Standard Hydrogen Electrode (SHE)

We cannot measure the potential of a single half-cell directly; we must measure the potential difference between two half-cells. To create a scale, we need a reference point, which is the Standard Hydrogen Electrode (SHE).

The SHE is defined as having an electrode potential of zero volts: \(E^\ominus = 0.00 \text{ V}\).

\(\text{2H}^{+}(\text{aq}) + 2\text{e}^- \rightarrow \text{H}_2(\text{g})\)       \(E^\ominus = 0.00 \text{ V}\)

Analogy: Think of the SHE as the "sea level" for measuring electrical potential. All other half-cell potentials are measured relative to this zero point.

The SHE setup requires:

• A platinum electrode (inert, provides a surface for the reaction).
• \(\text{H}^+\) ions at \(1.00 \text{ mol dm}^{-3}\) concentration.
• \(\text{H}_2\) gas passing over the electrode at \(100 \text{ kPa}\).
• Temperature maintained at \(298 \text{ K}\).

Predicting Redox Reactions using \(E^\ominus\) Values

When you look at a list of standard electrode potentials (the Electrochemical Series), the more positive the \(E^\ominus\) value, the greater the tendency for that species to be reduced.

The general rule for feasibility (i.e., whether a redox reaction will occur spontaneously) is:

The reaction will be feasible if the cell is constructed so that the reduction occurs in the half-cell with the MORE POSITIVE \(E^\ominus\) value.

The species on the left of the half-equation with the more positive \(E^\ominus\) will act as the oxidising agent (it gets reduced).
The species on the right of the half-equation with the more negative \(E^\ominus\) will act as the reducing agent (it gets oxidised).

Memory Aid: Positive Wins!

Imagine you have two half-cells, A and B.

• A has \(E^\ominus = +0.77 \text{ V}\)
• B has \(E^\ominus = -0.44 \text{ V}\)

Reaction A is the "winner" (more positive). It will proceed as written (reduction).
Reaction B is the "loser" (more negative). It will be forced to reverse (oxidation).

Calculating the Cell EMF (Electromotive Force)

The EMF (\(E_{\text{cell}}\)) is the maximum potential difference that the cell can generate. It is measured when virtually no current is drawn (this is why a high-resistance voltmeter is needed in Required Practical 6).

To calculate the standard EMF (\(E^\ominus_{\text{cell}}\)):

\(\mathbf{E^\ominus_{\text{cell}} = E^\ominus_{\text{reduction}} - E^\ominus_{\text{oxidation}}}\)

or, using the conventional cell notation (see below):

\(\mathbf{E^\ominus_{\text{cell}} = E^\ominus_{\text{RHS}} - E^\ominus_{\text{LHS}}}\)

If the calculated \(E^\ominus_{\text{cell}}\) is positive, the reaction is feasible (spontaneous).

Conventional Cell Representation (Cell Notation)

This is a shorthand way to write the structure of the electrochemical cell, starting with the oxidation side (anode) and ending with the reduction side (cathode).

Notation Rule:

\(\mathbf{\text{Reducing agent} \mid \text{Oxidised species} \parallel \text{Reduced species} \mid \text{Oxidising agent}}\)

Where:

• \(\mathbf{\mid}\) represents a phase boundary (e.g., solid metal into aqueous ions).
• \(\mathbf{\parallel}\) represents the Salt Bridge (allowing ion flow to maintain electrical neutrality).
• The left-hand side (LHS) is where Oxidation occurs.
• The right-hand side (RHS) is where Reduction occurs.

Example: A zinc-copper cell, where zinc is oxidised and copper is reduced.

\(\mathbf{\text{Zn}(\text{s}) \mid \text{Zn}^{2+}(\text{aq}, 1.0 \text{M}) \parallel \text{Cu}^{2+}(\text{aq}, 1.0 \text{M}) \mid \text{Cu}(\text{s})}\)

If inert electrodes (like Platinum, Pt) are needed (e.g., when no metal is present, like in a $\text{Fe}^{2+}/\text{Fe}^{3+}$ cell), the inert electrode is included at the end of the half-cell notation.
Example: \(\text{Pt}(\text{s}) \mid \text{Fe}^{2+}(\text{aq}), \text{Fe}^{3+}(\text{aq}) \parallel \ldots\)

3.1.9.2 Commercial Applications of Electrochemical Cells

Electrochemical cells are vital; they convert chemical energy directly into electrical energy. We use them as portable power sources. Cells are broadly classified by their reusability:

1. Non-rechargeable (Irreversible) Cells: Once the reactants are consumed, they cannot be regenerated. (e.g., standard alkaline batteries).
2. Rechargeable Cells: The chemical reactions are reversible. Applying an external voltage reverses the cell reaction, restoring the reactants. (e.g., car batteries, lithium cells).
3. Fuel Cells: Reactants (fuel and oxidant) are continuously supplied from an external source.

1. The Lithium-ion Cell (Rechargeable)

Lithium-ion cells power modern electronics due to their light weight and high energy density. The simplified reactions involve the movement of \(\text{Li}^+\) ions between the electrodes.

Charging/Discharging:

• During Discharging (Use), electrons flow from the negative electrode to the positive electrode.
• During Recharging, an external current forces electrons back the other way.

Simplified Electrode Reactions (Discharging Process):

• Negative Electrode (Anode, Oxidation): The lithium metal loses electrons.

  \(\mathbf{\text{Li} \rightarrow \text{Li}^+ + \text{e}^{-}}\)

• Positive Electrode (Cathode, Reduction): Lithium ions combine with cobalt oxide and electrons.

  \(\mathbf{\text{Li}^+ + \text{CoO}_2 + \text{e}^{-} \rightarrow \text{Li}^{+}[\text{CoO}_2]^{-}}\)

Did you know? The compound \(\text{Li}^{+}[\text{CoO}_2]^{-}\) formed at the positive electrode is an intercalation compound, meaning the \(\text{Li}^+\) ions are inserted between the layers of the cobalt oxide structure.

2. The Alkaline Hydrogen-Oxygen Fuel Cell

Fuel cells are different from batteries because they do not run down; they generate current as long as fuel is supplied. Hydrogen and oxygen are supplied continuously to the electrodes. The output is water, making it very environmentally friendly (zero greenhouse gas emissions).

Electrode Reactions (in Alkaline Conditions, e.g., using KOH electrolyte):

• Negative Electrode (Anode, Oxidation of Hydrogen): Hydrogen reacts with hydroxide ions to produce water and electrons.

  \(\mathbf{\text{2H}_2(\text{g}) + 4\text{OH}^{-}(\text{aq}) \rightarrow 4\text{H}_2\text{O}(\text{l}) + 4\text{e}^{-}}\)

• Positive Electrode (Cathode, Reduction of Oxygen): Oxygen gas gains electrons and reacts with water to form hydroxide ions.

  \(\mathbf{\text{O}_2(\text{g}) + 2\text{H}_2\text{O}(\text{l}) + 4\text{e}^{-} \rightarrow 4\text{OH}^{-}(\text{aq})}\)

Overall Reaction: Adding the two half-equations (and cancelling the \(\text{OH}^-\) and electrons) gives the simple overall reaction:

\(\mathbf{\text{2H}_2(\text{g}) + \text{O}_2(\text{g}) \rightarrow 2\text{H}_2\text{O}(\text{l})}\)

Benefits and Risks of Electrochemical Cells

Electrochemical cells (especially fuel cells and high-efficiency batteries) are crucial for developing sustainable energy.

Benefits:

• High Efficiency: Fuel cells convert chemical energy directly to electricity, avoiding the energy losses associated with combustion (heat loss).
• Environmentally Clean (Fuel Cells): They often produce only water as a byproduct.
• Portable Power: Batteries allow electrical devices to be portable and disconnected from the grid.

Risks/Disadvantages:

• Storage and Supply (Fuel Cells): Hydrogen fuel is often generated using fossil fuels, and storing hydrogen safely is challenging.
• Disposal (Non-rechargeable): Traditional batteries contain heavy metals (like Cadmium or Lead) which pose a disposal and environmental risk if not recycled properly.
• Rechargeable Cell Lifespan: Rechargeable cells degrade over time, limiting their usefulness and requiring eventual replacement.

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Key Takeaway Summary

Electrochemical cells allow us to harness redox reactions to create electricity.
The Standard Electrode Potential (\(E^\ominus\)) measures the tendency for reduction to occur under standard conditions, using the SHE (\(0.00 \text{ V}\)) as a reference point.
A spontaneous reaction occurs when the cell is set up so that the Reduction happens at the half-cell with the more positive \(E^\ominus\). The overall cell EMF must be positive.