Electrochemistry: Voltaic and Electrolytic Cells, Batteries, and Corrosion

Electrochemistry

Introduction to Electrochemistry

Electrochemistry is the study of chemical processes that involve the movement of electrons, resulting in the conversion between chemical and electrical energy. This field is fundamental to understanding batteries, fuel cells, corrosion, and industrial processes such as electrolysis.

• Electrochemical cells are devices that generate electrical energy from chemical reactions or use electrical energy to drive chemical reactions.

• There are two main types: voltaic (galvanic) cells and electrolytic cells.

Voltaic (Galvanic) Cells

Fuel Cells

Fuel cells are a type of voltaic cell that require a continuous supply of reactants (fuel and oxidant) to generate electricity. They are used in various applications, including vehicles and space missions.

• Anode reaction: $\mathrm{H_2(g) \rightarrow 2H^+(aq) + 2e^-}$

• Cathode reaction: $\mathrm{O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)}$

• Overall cell reaction: $\mathrm{2H_2(g) + O_2(g) \rightarrow 2H_2O(l)}$

• Fuel cells are efficient and produce water as the only byproduct, making them environmentally friendly.

Example: Hydrogen fuel cells are used in some cars and were originally developed for space applications.

Commercial Voltaic Cells

Several types of commercial voltaic cells are used in everyday devices. These include zinc-carbon dry cells, alkaline cells, lithium-iodine batteries, lead storage cells, and nickel-cadmium cells.

Zinc–Carbon Dry Cell (Leclanché Cell)

• Anode: $\mathrm{Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-}$

• Cathode: $\mathrm{2NH_4^+(aq) + 2MnO_2(s) + 2e^- \rightarrow Mn_2O_3(s) + H_2O(l) + 2NH_3(aq)}$

• Initial voltage is about 1.5 V but decreases rapidly in cold weather due to kinetic limitations.

Structure: The cell consists of a zinc can (anode), a graphite rod (cathode), and a paste of MnO2, ZnCl2, NH4Cl, and C.

Example 1

• The difference between a "heavy-duty" and a regular zinc-carbon battery is that the heavy-duty battery has a thicker zinc can, which increases its output and lifespan.

Zinc–Carbon Dry Cell: Alkaline

• Performs better than ordinary Zn-C batteries in cold weather.

• Anode (negative): $\mathrm{Zn(s) + 2OH^-(aq) \rightarrow Zn(OH)_2(s) + 2e^-}$

• Cathode (positive): $\mathrm{2MnO_2(s) + H_2O(l) + 2e^- \rightarrow Mn_2O_3(s) + 2OH^-(aq)}$

• Uses KOH as the electrolyte, which is an aqueous paste.

Problems with Zinc-Carbon Cells

• Side reactions between impurities and the electrolyte can increase self-discharge and promote corrosion.

• Not suitable for medical applications requiring long life and low current (e.g., pacemakers).

Lithium–Iodine Battery

• Solid-state battery with electrodes separated by a thin crystalline layer of lithium iodide.

• High resistance, low voltage at mA range, but very reliable and long-lasting (used in pacemakers).

Lead Storage Cell

• Electrodes are lead alloy grids: one with spongy lead (anode), one with lead dioxide (cathode), both in H2SO4 solution.

• Anode: $\mathrm{Pb(s) + HSO_4^-(aq) \rightarrow PbSO_4(s) + H^+(aq) + 2e^-}$

• Cathode: $\mathrm{PbO_2(s) + 3H^+(aq) + HSO_4^-(aq) + 2e^- \rightarrow PbSO_4(s) + 2H_2O(l)}$

• Each cell generates 2 V; six cells in series yield 12 V (typical car battery).

• Rechargeable; during recharge, reactions are reversed.

• Modern versions use calcium in the lead to reduce water decomposition, making them nearly maintenance-free.

Nickel–Cadmium Cell

• Anode: $\mathrm{Cd(s) + 2OH^-(aq) \rightarrow Cd(OH)_2(s) + 2e^-}$

• Cathode: $\mathrm{NiO(OH)(s) + H_2O(l) + e^- \rightarrow Ni(OH)_2(s) + OH^-(aq)}$

• Rechargeable; used in portable electronics.

• Nickel metal hydride and lithium hydride cells are less toxic alternatives.

Corrosion and Corrosion Control

Cathodic Protection

Corrosion is the deterioration of metals due to redox reactions with environmental agents. Cathodic protection is a method to prevent corrosion by making the metal to be protected the cathode of an electrochemical cell.

• Economic loss due to corrosion is significant, accounting for 2–4% of GDP in many countries.

• Most rusting occurs when water meets iron, forming a voltaic cell where iron is oxidized.

• Anode (iron): $\mathrm{Fe(s) \rightarrow Fe^{2+}(aq) + 2e^-}$

• Cathode (oxygen): $\mathrm{O_2(g) + 2H_2O(l) + 4e^- \rightarrow 4OH^-(aq)}$

• Connecting iron to a more active (more easily oxidized) metal like magnesium makes magnesium the anode and iron the cathode, protecting the iron from oxidation (sacrificial protection).

Example: Steel pipelines are protected from corrosion by attaching magnesium rods.

Experimental Illustration

• Iron corrosion yields Fe2+, which reacts with ferricyanide to give a dark blue precipitate.

• Where OH- forms, phenolphthalein turns pink.

Electrolytic Cells and Electrolysis

Electrolytic Cells

An electrolytic cell uses electrical energy to drive a nonspontaneous chemical reaction. This process is called electrolysis and is used industrially to produce substances like aluminum, chlorine, sodium, and magnesium.

Downs Cell

• Used for the electrolysis of molten sodium chloride to produce sodium metal and chlorine gas.

• Anode: $\mathrm{Cl^-(l) \rightarrow \frac{1}{2}Cl_2(g) + e^-}$

• Cathode: $\mathrm{Na^+(l) + e^- \rightarrow Na(l)}$

• Products must be kept separate to prevent recombination.

Historical Method

• Humphry Davy first produced sodium by electrolyzing molten sodium hydroxide ($318^\circ$C).

• Cathode: $\mathrm{Na^+(l) + e^- \rightarrow Na(l)}$

• Anode: $\mathrm{4OH^-(l) \rightarrow O_2(g) + 2H_2O(g) + 4e^-}$

Other Metals

• Lithium and magnesium are also obtained by electrolysis of their molten chlorides.

Aqueous Electrolysis

In aqueous solutions, water can also participate in the electrolysis reactions, so possible half-reactions must be considered for both the solute and water.

• Reduction (at cathode): $\mathrm{2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq)}$

• Oxidation (at anode): $\mathrm{2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-}$

• The reaction with the smallest voltage (least negative total potential) occurs.

Determining Electrolysis Reactions

1. Examine possible oxidation reactions; the one with the least negative $E^\circ$ (or most positive) occurs.

2. Examine possible reduction reactions; the one with the most positive $E^\circ$ occurs.

3. Select the reaction with the smallest negative voltage (opposite to galvanic cells, which maximize cell potential).

Example 2

• Electrolysis of sulfuric acid solution: Consider all possible half-cell reactions and select those with the smallest required voltage.

Overpotential

Overpotential is the difference between the theoretical (thermodynamic) reduction potential and the experimentally observed potential for a half-cell reaction. It arises due to factors such as electrode material, shape, and kinetic effects (e.g., gas bubble formation, charge transfer, diffusion, crystallization).

• Overpotential increases the voltage required for electrolytic cells and decreases the voltage available from galvanic cells.

• For most calculations, use standard reduction potentials unless otherwise specified.

Quantitative Aspects of Electrolysis

• Electric charge (Q): $Q = I \times t$

• Current (I): measured in amperes (A)

• Time (t): measured in seconds (s)

• Coulomb (C): $1\ \mathrm{C} = 1\ \mathrm{A} \cdot 1\ \mathrm{s}$

Applications and Examples

• Corrosion in Seawater vs. Freshwater: Seawater contains more ions, increasing the rate of corrosion compared to freshwater.

• Battery at Equilibrium: If all reactions in a battery reach equilibrium, no net current flows; turning on a device will not produce light or work.

• Redox Reaction Example: In a redox reaction between Cu and Al, Al is oxidized (loses electrons) and Cu is reduced (gains electrons) because Al is more active (more negative standard reduction potential).

Standard Reduction Potentials Table

The standard reduction potentials table lists half-cell reactions and their standard potentials (in volts) at 25°C. This table is essential for predicting the direction of redox reactions and calculating cell potentials.

Half-Reaction	Standard Potential (V)
Li+ + e- → Li(s)	-3.04
K+ + e- → K(s)	-2.93
Mg2+ + 2e- → Mg(s)	-2.37
Al3+ + 3e- → Al(s)	-1.66
Zn2+ + 2e- → Zn(s)	-0.76
Fe2+ + 2e- → Fe(s)	-0.44
Cu2+ + 2e- → Cu(s)	0.34
Ag+ + e- → Ag(s)	0.80
Cl2(g) + 2e- → 2Cl-	1.36
O2(g) + 4H+ + 4e- → 2H2O(l)	1.23
F2(g) + 2e- → 2F-	2.87

Additional info: Table entries are representative; refer to a full standard reduction potential table for more values.