Electrochemistry

Definition: Electrochemistry is the branch of chemistry that studies the relationship between electrical energy and chemical reactions, particularly redox reactions that produce or consume electric current.

How It Works

Every electrochemical cell separates a redox reaction into two half-reactions.
Oxidation, electron loss, happens at the anode; reduction, electron gain, happens at the cathode.
The two half-reactions are physically separated so electrons must travel through an external wire to get from one to the other.
That electron flow through the wire is the electric current.
Galvanic (voltaic) cells generate electricity spontaneously from a redox reaction that would happen anyway.
Electrons flow from anode to cathode through the external circuit.
Ions flow through a salt bridge or porous membrane to maintain charge neutrality in each half-cell.
In a galvanic cell, the anode is negative and the cathode is positive.
Electrolytic cells use an external power source to force a non-spontaneous redox reaction to occur.
Examples: splitting water into H2 and O2, or electroplating a thin metal layer onto an object.
Here the anode is positive and the cathode is negative, since the external source pushes electrons against their natural direction.
In both cell types, oxidation always happens at the anode and reduction always happens at the cathode.
That naming convention never flips, a useful mnemonic: “AnOx, RedCat.”
What flips between cell types is which electrode is positive versus negative.

Under the Hood

Standard reduction potentials (E°) are tabulated for half-reactions relative to the hydrogen electrode, defined as 0.00 V.

E°cell = E°cathode(reduction) - E°anode(reduction)

Worked example: Daniell cell.

Given:

Zn²+ + 2e- → Zn, E° = -0.76 V
Cu²+ + 2e- → Cu, E° = +0.34 V

Step 1: zinc is more easily oxidized (more negative reduction potential), so it’s the anode
Step 2: apply the cell potential formula

E°cell = E°cathode - E°anode = 0.34 - (-0.76)

Answer: E°cell = 1.10 V
Overall reaction:

Zn(s) + Cu²+(aq) → Zn²+(aq) + Cu(s)

Spontaneous as written, since E°cell is positive.

The relationship between cell potential and Gibbs free energy:

ΔG° = -nFE°

n is moles of electrons transferred, F is Faraday’s constant, 96,485 C/mol.
For the Daniell cell (n = 2):

ΔG° = -2 × 96,485 × 1.10

Answer: ΔG° ≈ -212 kJ/mol, confirming the reaction is strongly spontaneous.

The Nernst equation extends this to non-standard concentrations:

E = E° - (RT/nF)ln(Q)
E = E° - (0.0592/n)log(Q)  [at 25°C, base-10 log]

Worked example: Nernst equation.

Given: Daniell cell at 25°C, [Zn²+] = 0.010 M, [Cu²+] = 1.0 M
Step 1: find Q

Q = [Zn²+]/[Cu²+] = 0.010/1.0 = 0.010

Step 2: apply the Nernst equation

E = 1.10 - (0.0592/2)log(0.010) = 1.10 - (0.0296)(-2)

Answer: E = 1.159 V
Lowering the product-side ion concentration increases the cell potential slightly, consistent with Le Chatelier’s principle favoring the forward reaction.

Faraday’s laws of electrolysis quantify exactly how much product forms:

mass = (Q × M)/(n × F)

Q = current × time, in coulombs; M is molar mass; n is electrons per ion; F is Faraday’s constant.
This is the basis for calculating exact plating thickness and time in industrial electroplating.

Why It Matters

Electrochemistry powers batteries, from AA cells to lithium-ion packs.
It enables electroplating and corrosion-resistant coatings.
It drives large-scale industrial processes like the Hall-Héroult process for aluminum extraction.
That process consumes several percent of global electricity output.
Corrosion itself is an unwanted electrochemical process, iron oxidizing in the presence of water and oxygen.
Cathodic protection, attaching a more easily oxidized “sacrificial” metal like zinc or magnesium, prevents it.
That protection is a direct application of standard reduction potentials.

Common Pitfalls

Assuming the anode is always negative; that’s only true for galvanic cells.
In electrolytic cells the anode is positive because the external power source dictates the polarity.
Sign errors in the Nernst equation or in E°cell calculations.
A common one: forgetting to subtract the anode’s reduction potential rather than adding it.
Forgetting to balance electrons lost and gained between half-reactions before combining them.
Multiplying a half-reaction’s coefficients doesn’t change its E° value, E° is intensive, not extensive.
Confusing standard reduction potential tables, always written as reductions, with the actual direction a species reacts.
The less noble metal’s reaction must be reversed, and its E° sign flipped, when written as oxidation.
Mixing up amps, rate of charge flow, with coulombs, total charge, when calculating electroplating mass.
Faraday’s law calculations need total charge (current × time), not just current.
Forgetting that a more positive E°cell means a more spontaneous reaction, not “more concentrated” or “more reactive” in a general sense.

Comparison

Feature	Galvanic Cell	Electrolytic Cell
Reaction spontaneity	Spontaneous (ΔG < 0)	Non-spontaneous (ΔG > 0), forced
Energy	Produces electrical energy	Consumes electrical energy
Anode charge	Negative	Positive
Cathode charge	Positive	Negative
Example	AA battery, Daniell cell	Electroplating, water electrolysis

Quantity	Symbol	Relationship
Cell potential	E°	E°cell = E°cathode - E°anode
Free energy	ΔG°	ΔG° = -nFE°
Equilibrium constant	K	ΔG° = -RT ln K
Non-standard potential	E	E = E° - (RT/nF)ln Q
Mass deposited (electrolysis)	m	m = (Q × M)/(n × F)

Real-World Application

Cathodic protection guards buried pipelines and ship hulls from corrosion using electrochemistry directly.

A more easily oxidized metal, typically zinc or magnesium, is electrically connected to the steel structure.

Zn²+ + 2e- → Zn, E° = -0.76 V
Fe²+ + 2e- → Fe, E° = -0.44 V

Because zinc has a more negative standard reduction potential than iron, zinc preferentially oxidizes at the anode.
Zinc sacrifices itself instead of the steel.
Engineers size and replace these “sacrificial anodes” periodically.
This is essentially running a slow, deliberate galvanic cell to keep the far more expensive steel structure from corroding.

Example

A lithium-ion battery is a galvanic cell during discharge: lithium ions move from the graphite anode through an electrolyte to a metal-oxide cathode while electrons take the external path through the device being powered, releasing stored chemical energy as usable current. Charging the battery runs the same cell electrolytically, forcing lithium ions back to the anode using an external charger.

FAQ

Why do batteries eventually die even without being used?

Self-discharge: small internal side reactions and current leakage slowly convert reactants even with no external circuit connected.
This gradually reduces charge over weeks or months.

Can a single element have more than one standard reduction potential?

Yes, if it has multiple accessible oxidation states.
Iron has separate tabulated E° values for Fe³+ + e- → Fe²+ and Fe²+ + 2e- → Fe, different half-reactions.

Why is the hydrogen electrode used as the 0.00 V reference?

It’s an arbitrary but universally adopted convention.
All standard reduction potentials are really relative measurements.
Fixing one reaction at exactly 0.00 V gives every other half-reaction a consistent, comparable scale.

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