# Electrochemistry

Chemistry I · Energy, Equilibrium and Electrochemistry · https://tryals.app/learn/chemistry-i/electrochemistry

## Redox Separated in Space

Any redox reaction transfers electrons. If the two half-reactions are physically separated and connected by a wire, those electrons must travel through the circuit and can do electrical work. That is a **galvanic cell**.

**Oxidation** is loss of electrons and always happens at the **anode**; **reduction** is gain and always happens at the **cathode**. In a galvanic cell the anode is negative and the cathode positive. A **salt bridge** completes the circuit and keeps both compartments electrically neutral, without which the cell stops almost immediately.

Each half-reaction has a **standard electrode potential** $E^\circ$, measured against the standard hydrogen electrode defined as exactly 0 V. The cell potential is

$$E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$$

A positive $E^\circ_{cell}$ means a spontaneous reaction, tying electrochemistry to thermodynamics through

$$\Delta G^\circ = -nFE^\circ_{cell}$$

with $n$ the number of electrons transferred and $F = 96{,}485$ C/mol. Note that $E^\circ$ is an *intensive* property: it does not scale when you multiply a half-equation, even though $\Delta G^\circ$ does.

Away from standard conditions the **Nernst equation** applies:

$$E = E^\circ - \frac{0.0592}{n}\log Q$$

at 25 °C. As the reaction proceeds, $Q$ rises and $E$ falls, reaching zero exactly when $Q = K$, a flat battery is a cell that has reached equilibrium.

An **electrolytic cell** reverses the arrangement, using an external supply to drive a non-spontaneous reaction. The electrode names follow the chemistry, not the sign: oxidation is still at the anode, but in electrolysis the anode is positive.

> **Common pitfall:** multiplying a standard potential when balancing electrons. If a half-equation is doubled to balance a cell, its $\Delta G^\circ$ doubles but its $E^\circ$ does not change at all, potential is energy *per unit charge*, and both the energy and the charge scale together.

## Practice questions

7 of this lesson's 12 practice questions, with answers. The full set is in the app.

### 1. Why does a galvanic cell stop working almost immediately if the salt bridge is removed?

A. Charge builds up in each compartment and opposes further electron flow
B. The standard reduction potentials of the two half-cells become equal
C. Electrons can no longer complete their physical loop through the bridge
D. Both metal electrodes rapidly dissolve completely into their solutions

**Answer:** A. Charge builds up in each compartment and opposes further electron flow

**Why:** Oxidation leaves excess positive charge at the anode and reduction leaves excess negative charge at the cathode. That separation opposes further electron flow within moments unless ions migrate through a salt bridge to neutralise it.

Page: https://tryals.app/practice/chemistry-i/electrochemistry/why-does-a-galvanic-cell-stop-working-almost-immediately-if-the-salt

### 2. A cell has $E^\circ_{cathode} = +0.34$ V and $E^\circ_{anode} = -0.76$ V. Compute the standard cell potential in volts.

**Answer:** 1.1 (within ±0.02)

**Why:** $E^\circ_{cell} = 0.34 - (-0.76) = +1.10$ V. This is the Daniell cell of zinc and copper, and the positive value confirms the reaction is spontaneous as written.

Page: https://tryals.app/practice/chemistry-i/electrochemistry/a-cell-has-e-cathode-0-34-v-and-e-anode-0-76-v-compute-the

### 3. Electrode potential $E^\circ$ is an intensive property, whereas Gibbs energy $\Delta G^\circ$ is extensive. What follows from this distinction when balancing a complete redox reaction?

A. Multiplying a half-cell equation scales its $\Delta G^\circ$ but leaves $E^\circ$ unchanged
B. Multiplying a half-cell equation scales both $\Delta G^\circ$ and $E^\circ$ by the same factor
C. Multiplying a half-cell equation scales its $E^\circ$ but leaves $\Delta G^\circ$ unchanged
D. Multiplying a half-cell equation leaves both $\Delta G^\circ$ and $E^\circ$ strictly unchanged

**Answer:** A. Multiplying a half-cell equation scales its $\Delta G^\circ$ but leaves $E^\circ$ unchanged

**Why:** Thermodynamic potentials scale with the quantity of matter, but electrical potential is normalised per unit charge transferred. Confusing intensive voltages with extensive free energies leads to incorrect stoichiometric scaling during cell calculations.

Page: https://tryals.app/practice/chemistry-i/electrochemistry/electrode-potential-e-is-an-intensive-property-whereas-gibbs

### 4. Which statements about a galvanic cell are correct?

A. A flat battery is one whose reaction has run all the way to equilibrium
B. A positive cell potential indicates a spontaneous reaction
C. The standard potential doubles if the half-equation is doubled
D. Electrons flow through the external circuit from anode to cathode

**Answer:** A. A flat battery is one whose reaction has run all the way to equilibrium; B. A positive cell potential indicates a spontaneous reaction; D. Electrons flow through the external circuit from anode to cathode

**Why:** A positive potential means a negative $\Delta G$, electrons leave the oxidising anode, and the potential decays to zero as $Q$ climbs toward $K$. Potential is intensive and never scales with the equation.

Page: https://tryals.app/practice/chemistry-i/electrochemistry/which-statements-about-a-galvanic-cell-are-correct

### 5. Complete the description of the standard electrochemical reference points.

**Answer:** The **Faraday** constant is the charge carried by one mole of electrons, the standard **hydrogen** electrode is defined as exactly zero volts, and the **Nernst** equation gives the cell potential away from standard conditions.

**Why:** The Faraday constant converts moles of electrons into coulombs, the standard hydrogen electrode fixes the zero of the potential scale so every other potential is a comparison against it, and the Nernst equation handles non-standard concentrations.

Page: https://tryals.app/practice/chemistry-i/electrochemistry/complete-the-description-of-the-standard-electrochemical-reference

### 6. Arrange these metals in order of increasing tendency to be oxidised, given their standard reduction potentials: gold $+1.50$ V, copper $+0.34$ V, iron $-0.44$ V, zinc $-0.76$ V.

**Answer:**

1. Gold
2. Copper
3. Iron
4. Zinc

**Why:** The more negative the reduction potential, the more readily the metal is oxidised. Gold, at $+1.50$ V, resists oxidation almost entirely, while zinc at $-0.76$ V corrodes readily, which is exactly why zinc is used to protect steel.

Page: https://tryals.app/practice/chemistry-i/electrochemistry/arrange-these-metals-in-order-of-increasing-tendency-to-be-oxidised

### 7. Sort each feature by the type of cell it belongs to.

**Answer:**

- Galvanic cell: Converts chemical energy into electrical work, Has a positive cell potential
- Electrolytic cell: Requires an external power supply, Drives a non-spontaneous reaction, Has a positive anode

**Why:** A galvanic cell runs a spontaneous reaction and delivers work; an electrolytic cell consumes external power to force the reverse. Oxidation stays at the anode in both, but the electrode signs are opposite.

Page: https://tryals.app/practice/chemistry-i/electrochemistry/sort-each-feature-by-the-type-of-cell-it-belongs-to