Compare voltaic and electrolytic cells

Alberta Grade 12 Chemistry

Electrochemical Changes

Comparing Two Types of Electrochemical Cells — Alberta Chemistry 30, Outcome 30-B2.2k

Every time you use a battery-powered device, a chemical reaction is producing electricity. Every time you charge that same device, electricity is forcing a chemical reaction to happen. These two situations describe the two main types of electrochemical cells studied in Chemistry 30. A voltaic cell converts chemical energy into electrical energy. An electrolytic cell converts electrical energy into chemical energy. Both cells involve oxidation and reduction reactions (redox), but they differ in whether the reaction drives the current or the current drives the reaction. This lesson compares the two cell types in detail, including their components, electrode labels, energy direction, and the reactions taking place.

What you will learn

Identify the defining features of a voltaic cell and an electrolytic cell.
Compare the energy conversion, electrode labels, and spontaneity of reactions in each type of cell.
Predict the direction of electron flow and ion movement in each cell type.
Identify the anode and cathode in both cell types and state the reaction occurring at each electrode.

Prerequisite Review: Oxidation, Reduction, and Redox

Before comparing the two cell types, recall the core ideas from earlier in Chemistry 30. Oxidation is the loss of electrons by a species. Reduction is the gain of electrons by a species. The two processes always happen together, which is why they are called a redox reaction.

In any redox reaction, the species that loses electrons is called the reducing agent, and the species that gains electrons is called the oxidizing agent. Electrons released at one location must be accepted at another location. In an electrochemical cell, a physical path is set up so those electrons travel through a wire, creating a measurable electric current.

The Alberta Chemistry 30 Data Booklet contains a Standard Reduction Potential table. Each entry lists a half-reaction written as a reduction (gain of electrons). A species higher on the table has a greater tendency to be reduced. A species lower on the table has a greater tendency to be oxidized. This table is used to predict which reaction is spontaneous and to calculate cell voltage.

E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}

Oxidation = loss of electrons; reduction = gain of electrons.
Oxidation and reduction always occur together in a redox reaction.
The Standard Reduction Potential table ranks species by their tendency to be reduced.
A higher reduction potential means a stronger tendency to gain electrons.

The Voltaic Cell: Chemistry Produces Electricity

A voltaic cell (also called a galvanic cell) uses a spontaneous redox reaction to generate electrical energy. 'Spontaneous' means the reaction occurs on its own without being pushed by an outside energy source. A common observable example is placing a strip of zinc metal into a solution of copper(II) sulfate. The zinc slowly dissolves and a reddish layer of copper metal appears on the zinc surface. This happens because zinc is a stronger reducing agent than copper: zinc loses electrons to copper(II) ions without any outside help.

In a voltaic cell, this same chemistry is separated into two half-cells connected by a wire and a salt bridge. The half-cell where oxidation occurs is the anode. The half-cell where reduction occurs is the cathode. A helpful memory tool: the word 'oxidation' and 'anode' both start with vowels; the word 'reduction' and 'cathode' both start with consonants. Electrons flow through the external wire from the anode (where they are released) to the cathode (where they are consumed).

Because the reaction is spontaneous, the voltaic cell produces a positive cell voltage,

E^{\circ}_{\text{cell}} > 0

. The salt bridge allows ions to migrate between the two half-cells to maintain electrical neutrality, completing the internal circuit. In the anode half-cell, positive ions build up as the anode metal dissolves, so negative ions from the salt bridge migrate into that half-cell. In the cathode half-cell, positive ions are consumed as metal deposits, so positive ions from the salt bridge migrate in to replace them.

At the anode of a zinc–copper voltaic cell, the half-reaction is:

\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^{-}

. At the cathode, the half-reaction is:

\text{Cu}^{2+}(aq) + 2e^{-} \rightarrow \text{Cu}(s)

. The overall cell reaction is:

\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)

. Both atoms and charges balance in this equation.

E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}} > 0

A voltaic cell converts chemical energy into electrical energy.
The reaction is spontaneous: $E^{\circ}_{\text{cell}} > 0$ .
Anode = oxidation (negative terminal in a voltaic cell); cathode = reduction (positive terminal).
Electrons flow through the wire from anode to cathode.
The salt bridge maintains charge neutrality by allowing ion migration.

The Electrolytic Cell: Electricity Forces a Chemical Reaction

An electrolytic cell uses electrical energy supplied from an external source (such as a battery or power supply) to drive a non-spontaneous redox reaction. 'Non-spontaneous' means the reaction will not occur on its own; it needs a continuous input of energy to proceed. A familiar example is the electrolysis of water: applying a voltage across two electrodes submerged in water causes hydrogen gas to form at one electrode and oxygen gas at the other. Without the applied voltage, this decomposition does not happen.

Just as in a voltaic cell, oxidation happens at the anode and reduction happens at the cathode. This rule never changes regardless of cell type. In electrolysis, however, the anode is connected to the positive terminal of the external power supply and the cathode is connected to the negative terminal. Because the external source forces electrons to flow into the cathode, the anode in an electrolytic cell is considered the positive electrode and the cathode is the negative electrode — the opposite polarity from a voltaic cell.

Because the reaction is non-spontaneous, the calculated cell voltage using reduction potentials gives a negative value,

E^{\circ}_{\text{cell}} < 0

. The applied voltage from the external source must be at least equal in magnitude to overcome this negative value and make the reaction proceed. In practice, a slightly larger voltage (called the overpotential) is needed, but for Alberta Chemistry 30 comparisons the key idea is that an applied voltage is required.

A practical use of electrolytic cells is electroplating — depositing a thin layer of one metal onto another. For example, to chrome-plate a steel object, the steel object is made the cathode, a chromium electrode is the anode, and a chromium salt solution is the electrolyte. When current flows, chromium ions in solution are reduced and deposit as solid chromium on the steel object.

E^{\circ}_{\text{applied}} \geq |E^{\circ}_{\text{cell}}|

An electrolytic cell converts electrical energy into chemical energy.
The reaction is non-spontaneous: $E^{\circ}_{\text{cell}} < 0$ without the external source.
Anode = oxidation (positive electrode in an electrolytic cell); cathode = reduction (negative electrode).
Electrons are forced from cathode to the reaction site by an external power supply.
Applications include electroplating and industrial electrolysis.

Side-by-Side Comparison of the Two Cell Types

The most important skill for this outcome is being able to compare the two cell types clearly. Both types contain two electrodes, an electrolyte solution (or molten electrolyte), and a path for electron flow. The critical differences are the source of energy, the spontaneity of the reaction, the polarity of the electrodes, and the direction of energy conversion.

In both cell types, the anode is defined as the electrode where oxidation occurs and the cathode is defined as the electrode where reduction occurs. This definition is fixed. What changes between cell types is the sign (polarity) assigned to each electrode. In a voltaic cell, the anode carries a negative sign and the cathode carries a positive sign because the spontaneous reaction pushes electrons out of the anode. In an electrolytic cell, the external power supply forces electrons in at the cathode, making the cathode the negative electrode and the anode the positive electrode.

Ion movement in the electrolyte also deserves attention. In both cell types, cations (positive ions) migrate toward the cathode and anions (negative ions) migrate toward the anode. This is because the cathode environment becomes depleted in positive charge as cations are reduced, and the anode environment becomes enriched in positive charge as the anode is oxidized. The migration of ions through the solution completes the internal portion of the circuit.

\text{Oxidation at Anode} \quad \text{Reduction at Cathode}

Anode = oxidation and cathode = reduction in BOTH cell types — this never changes.
Voltaic cell: anode is negative, cathode is positive; electrolytic cell: anode is positive, cathode is negative.
Cations always migrate toward the cathode; anions always migrate toward the anode.
Voltaic: chemical energy → electrical energy; Electrolytic: electrical energy → chemical energy.

Calculating and Interpreting Cell Voltage

The standard cell voltage,

E^{\circ}_{\text{cell}}

, is calculated using reduction potentials from the Alberta Chemistry 30 Data Booklet. The formula is

E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}

, where both values are taken directly from the Standard Reduction Potential table as written (do not reverse the sign of either value before subtracting).

If

E^{\circ}_{\text{cell}}

is positive, the reaction is spontaneous as written and the cell operates as a voltaic cell. If

E^{\circ}_{\text{cell}}

is negative, the reaction is non-spontaneous as written and an external voltage must be applied — the cell must be set up as an electrolytic cell to make the reaction happen. This sign is the mathematical way to distinguish the two cell types.

The magnitude of

E^{\circ}_{\text{cell}}

tells you the driving force. A larger positive value means a stronger spontaneous tendency. A larger negative value means a larger external voltage must be supplied to force the reaction. Standard values assume 25°C, 1 mol/L concentrations, and 100 kPa pressure — the standard conditions stated in the Data Booklet.

E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}

$E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}$ using standard reduction potentials.
Positive $E^{\circ}_{\text{cell}}$ → spontaneous → voltaic cell.
Negative $E^{\circ}_{\text{cell}}$ → non-spontaneous → electrolytic cell setup required.
Standard conditions: 25°C, 1 mol/L, 100 kPa.

Voltaic Cell vs. Electrolytic Cell: Key Comparisons

Feature	Voltaic Cell	Electrolytic Cell
Energy conversion	Chemical → Electrical	Electrical → Chemical
Reaction spontaneity	Spontaneous	Non-spontaneous
Sign of $E^{\circ}_{\text{cell}}$	Positive ( $> 0$ )	Negative ( $< 0$ ) without applied voltage
External power supply?	No — cell produces current	Yes — required to drive reaction
Anode polarity	Negative electrode	Positive electrode
Cathode polarity	Positive electrode	Negative electrode
Reaction at anode	Oxidation	Oxidation
Reaction at cathode	Reduction	Reduction

Worked example

Identifying Cell Type and Calculating Cell Voltage

A cell is constructed using a nickel electrode in 1 mol/L

\text{Ni}^{2+}(aq)

and a silver electrode in 1 mol/L

\text{Ag}^{+}(aq)

, connected by a wire and a salt bridge. Using the Alberta Chemistry 30 Data Booklet, determine: (a) which electrode is the anode and which is the cathode, (b) the standard cell voltage

E^{\circ}_{\text{cell}}

, and (c) whether this cell operates as a voltaic or an electrolytic cell.

Step 1 — Look up standard reduction potentials.
Find the standard reduction potential for each half-reaction in the Data Booklet. Both values are written as reductions.
$\text{Ni}^{2+}(aq) + 2e^{-} \rightarrow \text{Ni}(s), \quad E^{\circ} = -0.26 \text{ V}$
Step 1 (continued) — Silver reduction potential.
The silver half-reaction has a higher (more positive) reduction potential, so silver ions have a greater tendency to be reduced.
$\text{Ag}^{+}(aq) + e^{-} \rightarrow \text{Ag}(s), \quad E^{\circ} = +0.80 \text{ V}$
Step 2 — Assign anode and cathode.
The species with the higher reduction potential is reduced at the cathode. The species with the lower reduction potential is oxidized at the anode. Silver has the higher value (+0.80 V), so the silver electrode is the cathode. Nickel has the lower value (−0.26 V), so the nickel electrode is the anode.
$\text{Cathode (reduction): } \text{Ag}^{+}(aq) + e^{-} \rightarrow \text{Ag}(s)$
Step 2 (continued) — Write the anode half-reaction.
At the anode, nickel is oxidized. The half-reaction is the reverse of the reduction listed in the Data Booklet.
$\text{Anode (oxidation): } \text{Ni}(s) \rightarrow \text{Ni}^{2+}(aq) + 2e^{-}$
Step 3 — Balance electrons and write the overall reaction.
The silver half-reaction involves 1 electron; the nickel half-reaction involves 2 electrons. Multiply the silver half-reaction by 2 so the electrons cancel, then add the half-reactions.
$\text{Ni}(s) + 2\text{Ag}^{+}(aq) \rightarrow \text{Ni}^{2+}(aq) + 2\text{Ag}(s)$
Step 4 — Calculate $E^{\circ}_{\text{cell}}$ .
Use the formula from the Data Booklet. Substitute the reduction potential of the cathode (Ag) and the reduction potential of the anode (Ni) directly — do not change the sign of either value before subtracting.
$E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}} = (+0.80 \text{ V}) - (-0.26 \text{ V}) = +1.06 \text{ V}$
Step 5 — Identify the cell type.
Because $E^{\circ}_{\text{cell}}$ is positive, the reaction is spontaneous. A spontaneous electrochemical cell is a voltaic cell. No external power supply is needed.
$E^{\circ}_{\text{cell}} = +1.06 \text{ V} > 0 \Rightarrow \text{voltaic cell}$

Answer: The nickel electrode is the anode (oxidation) and the silver electrode is the cathode (reduction). The standard cell voltage is

E^{\circ}_{\text{cell}} = +1.06 \text{ V}

. Because

E^{\circ}_{\text{cell}} > 0

, this is a voltaic cell — the reaction is spontaneous.

Check: Verify atom balance: 1 Ni on each side ✓; 2 Ag on each side ✓. Verify charge balance: left side has

2 \times (+1) = +2

; right side has

+2

from

\text{Ni}^{2+}

✓. The sign of

E^{\circ}_{\text{cell}}

is positive, confirming spontaneity and voltaic operation ✓.

Common mistakes and how to avoid them

Reversing the sign of the reduction potential before using the cell voltage formula — for example, changing

E^{\circ}_{\text{anode}}

to a positive value before subtracting.

Correction: Use both reduction potentials exactly as listed in the Data Booklet. The formula

E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}

already accounts for the direction change at the anode. Do not flip the sign of the anode value separately.

Thinking the anode is always positive or the cathode is always negative, regardless of cell type.

Correction: In a voltaic cell the anode is negative and the cathode is positive. In an electrolytic cell the anode is positive and the cathode is negative. The electrode labels (anode/cathode) are fixed to the reaction type (oxidation/reduction), but the polarity depends on whether the cell is voltaic or electrolytic.

Forgetting to balance the number of electrons transferred when combining half-reactions, leading to an unbalanced overall equation.

Correction: Always check that the electrons released at the anode equal the electrons consumed at the cathode before adding the half-reactions. Multiply one or both half-reactions by a whole number if needed.

Confusing the direction of ion migration — for example, stating that cations move toward the anode.

Correction: Cations (positive ions) always migrate toward the cathode because the cathode attracts positive species. Anions (negative ions) always migrate toward the anode. This rule applies to both voltaic and electrolytic cells.

Concluding that a cell is electrolytic simply because it contains a metal electrode or liquid electrolyte, rather than checking

E^{\circ}_{\text{cell}}

.

Correction: The correct way to identify the cell type is to calculate

E^{\circ}_{\text{cell}}

. A positive value means voltaic (spontaneous); a negative value means the cell can only function with an applied external voltage, making it electrolytic.

Lesson summary

A voltaic cell converts chemical energy into electrical energy through a spontaneous redox reaction, giving $E^{\circ}_{\text{cell}} > 0$ .
An electrolytic cell converts electrical energy into chemical energy by forcing a non-spontaneous redox reaction using an external power supply, giving $E^{\circ}_{\text{cell}} < 0$ without that supply.
In both cell types, oxidation occurs at the anode and reduction occurs at the cathode — this definition never changes.
Electrode polarity differs: in a voltaic cell the anode is negative and the cathode is positive; in an electrolytic cell the anode is positive and the cathode is negative.
Cations migrate toward the cathode and anions migrate toward the anode in both cell types, completing the internal circuit.
Cell voltage is calculated as $E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}$ using standard reduction potentials from the Alberta Chemistry 30 Data Booklet.

Check your understanding

Question 1

A cell is built using a copper electrode and a zinc electrode. The calculated

E^{\circ}_{\text{cell}}

is

+1.10

V. Which statement correctly describes this cell?

It is an electrolytic cell because copper is reduced at the cathode.
It is a voltaic cell because the positive cell voltage indicates a spontaneous reaction.
It is an electrolytic cell because an external power supply is needed to produce +1.10 V.
It is a voltaic cell because zinc is always the cathode in any electrochemical cell.

Show answer and explanation

It is a voltaic cell because the positive cell voltage indicates a spontaneous reaction.

A positive

E^{\circ}_{\text{cell}}

means the reaction is spontaneous. A spontaneous electrochemical cell is a voltaic cell — it produces electrical energy on its own. Option A is wrong because the cell type is not determined by which ion is reduced. Option C is wrong because the positive sign shows no external supply is needed. Option D is wrong because electrode identity depends on reduction potentials, not a fixed rule for a given metal.

Question 2

In any electrochemical cell, which statement about the anode is always true?

The anode is the positive electrode.
The anode is where reduction takes place.
The anode is where oxidation takes place.
The anode receives electrons from the external circuit.

Show answer and explanation

The anode is where oxidation takes place.

By definition, the anode is always the electrode where oxidation (loss of electrons) takes place. This is true for both voltaic and electrolytic cells. The polarity of the anode (positive or negative) changes depending on cell type, so option A is only sometimes true. Option B describes the cathode. Option D describes the cathode, because the cathode receives electrons through the external wire.

Question 3

During the electrolysis of a molten salt, which direction do cations move through the electrolyte?

Toward the anode, because the anode is positive in an electrolytic cell.
Toward the cathode, because cations are attracted to the negative electrode.
Randomly, because there is no net charge difference between the electrodes.
Toward the anode, because oxidation releases electrons that attract cations.

Show answer and explanation

Toward the cathode, because cations are attracted to the negative electrode.

Cations carry a positive charge and are attracted to the cathode, which is the negative electrode in an electrolytic cell. At the cathode, cations gain electrons and are reduced. Option A incorrectly reverses the direction. Option C is wrong because there is a definite charge difference between the electrodes. Option D is wrong because the electrons released at the anode leave through the external wire, not into the solution.

Question 4

The standard reduction potentials for two half-reactions are:

X^{2+}(aq) + 2e^{-} \rightarrow X(s)

,

E^{\circ} = -0.44

V, and

Y^{2+}(aq) + 2e^{-} \rightarrow Y(s)

,

E^{\circ} = +0.34

V. If a cell is constructed with X as the anode and Y as the cathode, what is

E^{\circ}_{\text{cell}}

and what type of cell is it?

$E^{\circ}_{\text{cell}} = -0.78$ V; electrolytic cell
$E^{\circ}_{\text{cell}} = +0.78$ V; voltaic cell
$E^{\circ}_{\text{cell}} = -0.10$ V; electrolytic cell
$E^{\circ}_{\text{cell}} = +0.10$ V; voltaic cell

Show answer and explanation

E^{\circ}_{\text{cell}} = +0.78

V; voltaic cell

Using the formula

E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}

:

E^{\circ}_{\text{cell}} = (+0.34 \text{ V}) - (-0.44 \text{ V}) = +0.78 \text{ V}

. The positive result confirms the reaction is spontaneous, so this is a voltaic cell. Option A incorrectly subtracts in the wrong order. Options C and D use the wrong arithmetic.

Key terms

Voltaic cell: An electrochemical cell in which a spontaneous redox reaction converts chemical energy into electrical energy. Also called a galvanic cell.
Electrolytic cell: An electrochemical cell in which electrical energy from an external source drives a non-spontaneous redox reaction, converting electrical energy into chemical energy.
Anode: The electrode where oxidation (loss of electrons) occurs in any electrochemical cell.
Cathode: The electrode where reduction (gain of electrons) occurs in any electrochemical cell.
Standard reduction potential ( $E^{\circ}$ ): A measure of the tendency of a species to gain electrons under standard conditions (25°C, 1 mol/L, 100 kPa), listed in the Alberta Chemistry 30 Data Booklet.
Standard cell voltage ( $E^{\circ}_{\text{cell}}$ ): The overall voltage produced (or required) by an electrochemical cell under standard conditions, calculated as $E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}$ .
Spontaneous reaction: A reaction that proceeds on its own without a continuous input of energy from an external source. In electrochemistry, indicated by $E^{\circ}_{\text{cell}} > 0$ .
Electrolysis: The process of using an applied electric current to drive a non-spontaneous chemical reaction in an electrolytic cell.

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Published by DoAssignment. This reviewed lesson follows Alberta Grade 12 Chemistry (CHEM30), expectation 30-B2.2k. It is a study resource, not an official curriculum publication.

Before publication, content is checked for structure, mathematical or chemical notation, calculations, course boundaries, and readability. Errors can still occur, so corrections are welcomed.

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30-B2.2k · Compare voltaic and electrolytic cells

Alberta Grade 12 Chemistry

Comparing Two Types of Electrochemical Cells — Alberta Chemistry 30, Outcome 30-B2.2k

What you will learn

Prerequisite Review: Oxidation, Reduction, and Redox

The Voltaic Cell: Chemistry Produces Electricity

The Electrolytic Cell: Electricity Forces a Chemical Reaction

Side-by-Side Comparison of the Two Cell Types

Calculating and Interpreting Cell Voltage

Voltaic Cell vs. Electrolytic Cell: Key Comparisons

Identifying Cell Type and Calculating Cell Voltage

Common mistakes and how to avoid them

Lesson summary

Check your understanding

Question 1

Question 2

Question 3

Question 4

Key terms

Continue through CHEM30

About this lesson and its review