Electrochemical cells have massive commercial value. They power everything from small household devices and mobile phones to electric cars and spacecraft. Depending on their chemistry and design, commercial cells are classified as non-rechargeable (primary), rechargeable (secondary), or fuel cells.
A cell in which the chemical reactions that generate electricity are irreversible. Once the reactants are completely converted to products, the cell potential falls to zero and it must be discarded.
A cell in which the chemical reactions that generate electricity are reversible. By applying an external electrical voltage, the cell reactions are driven backwards, reforming the original reactants so the cell can be used again.
An electrochemical cell that generates electricity by continuously reacting an external fuel supply (typically hydrogen) with an oxidant (typically oxygen) in the presence of an electrolyte.
Rechargeable Cells: The Lithium-Ion Battery
Lithium-ion cells are the most common rechargeable battery in modern electronics due to their high energy density and light weight. They consist of a negative carbon (graphite) electrode containing intercalated lithium atoms, a positive lithium cobalt oxide (\( \text{LiCoO}_2 \)) electrode, and an organic solvent electrolyte containing dissolved lithium salts.
🔑 Key Principle
During discharging (powering a device), lithium atoms in the graphite host release electrons to form lithium ions. The ions migrate through the internal electrolyte to the positive electrode, while the electrons travel through the external circuit to do work. Charging reverses this process using an external voltage.
The half-equations during discharging are:
- Negative Electrode (Oxidation):
\( \text{LiC}_6\text{(s)} \rightarrow \text{Li}^+\text{(aq)} + \text{C}_6\text{(s)} + \text{e}^- \)
- Positive Electrode (Reduction):
\( \text{Li}^+\text{(aq)} + \text{CoO}_2\text{(s)} + \text{e}^- \rightarrow \text{LiCoO}_2\text{(s)} \)
- Overall Discharge Equation:
\( \text{LiC}_6\text{(s)} + \text{CoO}_2\text{(s)} \rightarrow \text{LiCoO}_2\text{(s)} + \text{C}_6\text{(s)} \)
Fuel Cells: The Hydrogen-Oxygen Fuel Cell
Unlike conventional batteries which store their reactants internally, a fuel cell is fed reactants continuously from an external tank. The most common type is the hydrogen-oxygen fuel cell. It converts the chemical energy of hydrogen and oxygen directly into electrical energy, producing only water as a chemical waste product.
Fuel Cell Chemistry by Electrolyte
The electrochemical reactions inside a hydrogen-oxygen fuel cell depend on whether the electrolyte is alkaline or acidic. You must know both sets of half-equations for AQA exams.
A classic exam trap is combining alkaline and acidic half-equations or forgetting to cancel spectator water molecules. Notice that in alkaline conditions, hydroxide ions (\( \text{OH}^- \)) appear, while in acidic conditions, hydrogen ions (\( \text{H}^+ \)) appear. No matter which electrolyte is used, when you add the anode and cathode half-equations, the overall balanced cell equation is exactly the same!
Case 1: Alkaline Electrolyte (e.g. Potassium Hydroxide, KOH)
Hydroxide ions flow through the electrolyte to complete the circuit:
- Anode (Oxidation):
\( \text{H}_2\text{(g)} + 2\text{OH}^-\text{(aq)} \rightarrow 2\text{H}_2\text{O(l)} + 2\text{e}^- \)
- Cathode (Reduction):
\( \text{O}_2\text{(g)} + 2\text{H}_2\text{O(l)} + 4\text{e}^- \rightarrow 4\text{OH}^-\text{(aq)} \)
Case 2: Acidic Electrolyte (e.g. Phosphoric Acid or Proton Exchange Membrane)
Hydrogen ions (protons) flow through the membrane/electrolyte:
- Anode (Oxidation):
\( \text{H}_2\text{(g)} \rightarrow 2\text{H}^+\text{(aq)} + 2\text{e}^- \)
- Cathode (Reduction):
\( \text{O}_2\text{(g)} + 4\text{H}^+\text{(aq)} + 4\text{e}^- \rightarrow 2\text{H}_2\text{O(l)} \)
Overall Equation (Alkaline & Acidic)
\( 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)} \quad E^\theta_{\text{cell}} = +1.23\text{ V} \)
Pros and Cons of Hydrogen Fuel Cells
Hydrogen fuel cells are heavily researched as a clean alternative to the internal combustion engine. However, they face several real-world challenges:
| Advantages | Disadvantages |
|---|---|
| Zero direct emissions: The only chemical product is water. No carbon dioxide, carbon monoxide, or nitrogen oxides are released. | Production footprints: Most commercial hydrogen is currently made by steam reforming methane, a process that releases carbon dioxide. |
| High efficiency: They convert chemical energy directly to electricity, avoiding the thermal energy losses of combustion engines. | Storage and transport: Hydrogen is a highly flammable gas with low energy density by volume. It must be compressed to high pressures or liquefied. |
| Constant power: Unlike batteries, they do not need recharging and will run indefinitely as long as fuel is supplied. | High cost: The electrodes require expensive catalysts, typically platinum, making the cells expensive to manufacture. |
Step 1: Write down the overall discharge equation:
\( \text{LiC}_6\text{(s)} + \text{CoO}_2\text{(s)} \rightarrow \text{LiCoO}_2\text{(s)} + \text{C}_6\text{(s)} \)
Step 2: Reverse the equation for charging:
During charging, we apply an external voltage to force the reverse process to occur:
\( \text{LiCoO}_2\text{(s)} + \text{C}_6\text{(s)} \rightarrow \text{LiC}_6\text{(s)} + \text{CoO}_2\text{(s)} \)
Step 3: State the movement of lithium ions:
During charging, lithium ions (\( \text{Li}^+ \)) are released from the positive lithium cobalt oxide (\( \text{LiCoO}_2 \)) electrode, travel through the electrolyte, and are inserted (intercalated) back into the negative graphite (\( \text{C}_6 \)) electrode.
Step 1: Add the alkaline half-equations:
Anode (oxidation): \( 2\text{H}_2\text{(g)} + 4\text{OH}^-\text{(aq)} \rightarrow 4\text{H}_2\text{O(l)} + 4\text{e}^- \) (multiplied by 2 to balance electrons)
Cathode (reduction): \( \text{O}_2\text{(g)} + 2\text{H}_2\text{O(l)} + 4\text{e}^- \rightarrow 4\text{OH}^-\text{(aq)} \)
Add them together:
\( 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} + 4\text{OH}^-\text{(aq)} + 2\text{H}_2\text{O(l)} + 4\text{e}^- \rightarrow 4\text{H}_2\text{O(l)} + 4\text{OH}^-\text{(aq)} + 4\text{e}^- \)
Cancel out the \( 4\text{e}^- \), \( 4\text{OH}^- \), and subtract \( 2\text{H}_2\text{O} \) from both sides:
\( 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)} \)
Step 2: Add the acidic half-equations:
Anode (oxidation): \( 2\text{H}_2\text{(g)} \rightarrow 4\text{H}^+\text{(aq)} + 4\text{e}^- \) (multiplied by 2 to balance electrons)
Cathode (reduction): \( \text{O}_2\text{(g)} + 4\text{H}^+\text{(aq)} + 4\text{e}^- \rightarrow 2\text{H}_2\text{O(l)} \)
Add them together:
\( 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} + 4\text{H}^+\text{(aq)} + 4\text{e}^- \rightarrow 2\text{H}_2\text{O(l)} + 4\text{H}^+\text{(aq)} + 4\text{e}^- \)
Cancel out the \( 4\text{e}^- \) and \( 4\text{H}^+ \):
\( 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)} \)
Conclusion: Both sets of half-equations sum to the identical overall redox process.
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