Predict redox spontaneity using standard reduction potentials

Alberta Grade 12 Chemistry

Electrochemical Changes

Alberta Chemistry 30 — Outcome 30-B1.6k

Drop a small piece of zinc metal into a blue copper(II) sulfate solution and you will see the blue colour fade, a reddish-brown solid appear on the zinc, and the zinc slowly dissolve. Nothing needs to push this reaction — it happens on its own. In contrast, dropping a copper strip into a colourless zinc sulfate solution produces no visible change at all. One direction goes, the other does not. Why? The answer comes from comparing the tendencies of different species to gain electrons. This lesson shows you how to use the numbers in the Alberta Chemistry 30 Data Booklet — called standard reduction potentials — to predict which direction a redox reaction will proceed spontaneously, before ever running an experiment.

What you will learn

Explain what a standard reduction potential measures and how to read its value from the Alberta Chemistry 30 Data Booklet table.
Identify the oxidizing agent and the reducing agent from a pair of half-reactions.
Calculate the standard cell potential for a redox reaction using the reduction potentials of the two half-reactions.
Use the sign of the standard cell potential to predict whether a redox reaction is spontaneous under standard conditions.
Apply the spontaneity prediction to multi-step or unfamiliar redox pairs at mixed difficulty levels.

Prerequisite Review: Oxidation, Reduction, and Half-Reactions

A redox reaction always involves two simultaneous processes. Oxidation is the loss of electrons by a species; reduction is the gain of electrons by a species. The species that loses electrons is called the reducing agent, and the species that gains electrons is called the oxidizing agent. These ideas come from Chemistry 20 and are the foundation for everything in this lesson.

Every redox reaction can be split into two half-reactions. The oxidation half-reaction shows electrons being released, and the reduction half-reaction shows electrons being absorbed. For example, the zinc–copper reaction separates into two half-reactions: zinc loses two electrons, and copper(II) ions each gain two electrons. Both half-reactions must be balanced for atoms and for charge before they can be combined.

A key rule: when you add the two half-reactions together to form the overall equation, the electrons released in oxidation must exactly equal the electrons consumed in reduction. If the numbers do not match, you must multiply one or both half-reactions by a whole number so that they cancel out completely. This is the same electron-balancing technique used throughout Chemistry 30 electrochemistry.

Oxidation = loss of electrons; reduction = gain of electrons.
The reducing agent is oxidized; the oxidizing agent is reduced.
Half-reactions show electrons explicitly and must balance atoms and charge.
Electrons must cancel completely when half-reactions are combined.

Standard Reduction Potentials: What They Mean

Every half-reaction has a measurable tendency to proceed as a reduction — that is, a tendency to pull electrons toward itself. This tendency is assigned a number called the standard reduction potential, given the symbol E°. The unit is the volt (V). The word 'standard' means the measurement was made with all dissolved species at 1 mol/L concentration, all gases at 101.3 kPa pressure, and at 25°C.

The Alberta Chemistry 30 Data Booklet lists half-reactions in the form: oxidized form + electrons → reduced form. Each half-reaction is written as a reduction, and its E° value sits beside it. A large positive E° means the species strongly wants to gain electrons — it is a strong oxidizing agent. A large negative E° means the species strongly resists gaining electrons — its reverse reaction (oxidation) is more favourable, making it a good reducing agent.

The reference point of the entire table is the standard hydrogen electrode (SHE):

2\,\text{H}^+(aq) + 2\,e^- \rightarrow \text{H}_2(g)

, which is assigned exactly

E° = 0.00\text{ V}

. Every other value in the table is measured relative to this reference. Half-reactions listed above the SHE in the Data Booklet have positive E° values and are stronger oxidizing agents than

\text{H}^+

. Those listed below have negative E° values and are weaker oxidizing agents than

\text{H}^+

.

An important rule: the E° value of a half-reaction does not change when you multiply the half-reaction by a whole number. Cell potentials are intensive quantities — they do not depend on how many moles react. This is different from enthalpy changes, where multiplying a reaction does scale the value. Keep this distinction clear to avoid a very common error.

E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}

E° is the standard reduction potential, measured in volts (V), listed in the Data Booklet for each half-reaction written as a reduction.
A more positive E° indicates a stronger tendency to be reduced (stronger oxidizing agent).
A more negative E° indicates a stronger tendency to be oxidized (stronger reducing agent).
The SHE ( $E° = 0.00\text{ V}$ ) is the reference for all values.
Multiplying a half-reaction by a whole number does not change its E° value.

Calculating the Standard Cell Potential

To predict spontaneity, you calculate the standard cell potential,

E°_{\text{cell}}

. The cathode is the electrode where reduction occurs; the anode is where oxidation occurs. The Alberta Chemistry 30 Data Booklet formula for the standard cell potential is:

E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}

. Both

E°_{\text{cathode}}

and

E°_{\text{anode}}

are looked up directly from the Data Booklet table as reduction potentials — you subtract the anode's reduction potential from the cathode's reduction potential. You do not flip the sign of the anode value manually; the subtraction does that for you.

The rule for spontaneity is direct: if

E°_{\text{cell}}

is positive, the reaction is spontaneous in the forward direction under standard conditions. If

E°_{\text{cell}}

is negative, the reaction is non-spontaneous as written — the reverse reaction would be spontaneous instead. If

E°_{\text{cell}}

equals zero, the system is at equilibrium and neither direction is favoured.

To identify which species is reduced and which is oxidized, compare the two reduction potentials from the Data Booklet. The half-reaction with the higher (more positive) E° will proceed as written — as a reduction — and becomes the cathode. The half-reaction with the lower (more negative) E° is reversed — it proceeds as an oxidation — and becomes the anode. The species with the higher reduction potential is therefore the oxidizing agent, and the species with the lower reduction potential is the reducing agent.

Always keep track of your labels. Write down which half-reaction you assign to the cathode and which to the anode before calculating. This organized approach prevents sign errors, which are the most frequent source of mistakes on diploma-style questions.

E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}

Cathode = reduction; anode = oxidation.
$E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}$ (both values taken as reduction potentials from the Data Booklet).
Positive $E°_{\text{cell}}$ → spontaneous forward reaction.
Negative $E°_{\text{cell}}$ → non-spontaneous forward reaction (reverse is spontaneous).
The half-reaction with the higher E° is always the cathode (reduction).

Connecting the Particle Model to the Numbers

At the particle level, a spontaneous redox reaction occurs because one species pulls electrons away from another more strongly than that second species resists losing them. In the zinc–copper example, copper(II) ions have a much stronger pull on electrons (

E° = +0.34\text{ V}

) than zinc ions do (

E° = -0.76\text{ V}

). When zinc atoms and copper(II) ions meet, electrons flow from zinc to copper(II) ions. Zinc atoms are oxidized at the particle level: each zinc atom gives up two electrons to a nearby copper(II) ion, which is then neutralized and deposits as copper metal.

This electron transfer is the particle-level explanation for every observation in the opening phenomenon: the blue colour fades because

\text{Cu}^{2+}

ions (which give the solution its colour) are removed from solution; the reddish-brown solid appears because neutral copper atoms are deposited; and the zinc dissolves because zinc atoms enter solution as

\text{Zn}^{2+}

ions.

The standard cell potential

E°_{\text{cell}}

is essentially a measure of how large that electron-pulling difference is. A larger positive value means a larger 'pull' difference and a more strongly spontaneous reaction. A value near zero means the two species have nearly equal electron-pulling tendencies, and the reaction barely favours one direction.

Spontaneous electron transfer flows from the species with lower E° to the species with higher E°.
A larger positive $E°_{\text{cell}}$ corresponds to a greater difference in electron-attracting ability and a more strongly spontaneous reaction.
Observable changes (colour, solid formation, dissolving) map directly onto which species gains and which loses electrons at the particle level.

Applying the Method: Step-by-Step Strategy

Use this four-step strategy for any redox spontaneity question. Step 1: Identify the two half-reactions involved and look up their standard reduction potentials in the Alberta Chemistry 30 Data Booklet. Step 2: Assign the cathode (higher E°, reduction) and the anode (lower E°, oxidation). Step 3: Apply the formula

E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}

using the reduction potential values directly from the table. Step 4: State your conclusion — if

E°_{\text{cell}} > 0

, the reaction is spontaneous as written; if

E°_{\text{cell}} < 0

, it is not spontaneous as written.

For harder questions, you may be given the overall redox equation and asked to predict spontaneity without being told which species is oxidized or reduced. In those cases, assign oxidation numbers to identify what changed, then match the changes to the correct half-reactions in the Data Booklet. The four-step strategy still applies once you have identified the half-reactions.

Remember: never change the E° value from the table when you scale a half-reaction. If the electron counts in the two half-reactions do not match, multiply the half-reactions by appropriate integers, but leave the E° values unchanged. This is one of the most tested points on the Alberta Chemistry 30 diploma examination.

E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}

Four steps: (1) look up E° values, (2) assign cathode and anode, (3) calculate $E°_{\text{cell}}$ , (4) state spontaneity.
For harder questions, use oxidation numbers to identify which species is oxidized and which is reduced before looking up half-reactions.
Never multiply E° values when scaling half-reactions.

Standard Reduction Potentials for Selected Half-Reactions (from Alberta Chemistry 30 Data Booklet)

Half-Reaction (written as reduction)	E° (V)
$\text{F}_2(g) + 2\,e^- \rightarrow 2\,\text{F}^-(aq)$	+2.87
$\text{Cl}_2(g) + 2\,e^- \rightarrow 2\,\text{Cl}^-(aq)$	+1.36
$\text{Cu}^{2+}(aq) + 2\,e^- \rightarrow \text{Cu}(s)$	+0.34
$2\,\text{H}^+(aq) + 2\,e^- \rightarrow \text{H}_2(g)$	0.00
$\text{Zn}^{2+}(aq) + 2\,e^- \rightarrow \text{Zn}(s)$	−0.76
$\text{Mg}^{2+}(aq) + 2\,e^- \rightarrow \text{Mg}(s)$	−2.37
$\text{Li}^+(aq) + e^- \rightarrow \text{Li}(s)$	−3.04

Worked example

Predicting Spontaneity for the Zinc–Copper Reaction

Predict whether the following reaction is spontaneous under standard conditions. Identify the oxidizing agent and the reducing agent. Show your calculation using the Alberta Chemistry 30 Data Booklet values.

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

Step 1 — Identify the two half-reactions and look up their standard reduction potentials.
Zinc is oxidized (it loses electrons, going from 0 to +2). Copper(II) is reduced (it gains electrons, going from +2 to 0). Look up both half-reactions in the Alberta Chemistry 30 Data Booklet, where they appear as reductions.
$\text{Cu}^{2+}(aq) + 2\,e^- \rightarrow \text{Cu}(s) \quad E° = +0.34\text{ V}$
Step 1 continued — Zinc reduction half-reaction from the Data Booklet.
The zinc half-reaction as a reduction is listed in the Data Booklet. Its standard reduction potential is negative, meaning zinc ions do not strongly attract electrons.
$\text{Zn}^{2+}(aq) + 2\,e^- \rightarrow \text{Zn}(s) \quad E° = -0.76\text{ V}$
Step 2 — Assign cathode and anode.
The half-reaction with the higher (more positive) reduction potential is the cathode — reduction occurs here. The half-reaction with the lower reduction potential is the anode — oxidation occurs here. Compare +0.34 V and −0.76 V: copper(II) has the higher value, so it is reduced at the cathode. Zinc is oxidized at the anode.
$E°_{\text{cathode}} = +0.34\text{ V} \quad (\text{Cu}^{2+}/\text{Cu}) \qquad E°_{\text{anode}} = -0.76\text{ V} \quad (\text{Zn}^{2+}/\text{Zn})$
Step 3 — Calculate the standard cell potential.
Apply the Data Booklet formula. Both values are taken directly from the table as reduction potentials. The subtraction automatically accounts for the reversal of the anode half-reaction.
$E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}} = (+0.34\text{ V}) - (-0.76\text{ V}) = +1.10\text{ V}$
Step 4 — State the conclusion about spontaneity and identify the agents.
Because the standard cell potential is positive (+1.10 V > 0), the reaction is spontaneous as written under standard conditions. The species that is reduced (gains electrons) is the oxidizing agent; the species that is oxidized (loses electrons) is the reducing agent.
$E°_{\text{cell}} = +1.10\text{ V} > 0 \implies \text{spontaneous}$

Answer: The reaction is spontaneous under standard conditions.

E°_{\text{cell}} = +1.10\text{ V}

. The oxidizing agent is

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

(it is reduced). The reducing agent is

\text{Zn}(s)

(it is oxidized).

Check: Verify by checking direction: zinc (

E° = -0.76\text{ V}

) has the lower reduction potential, so it is correctly assigned as the anode (oxidation). Copper(II) (

E° = +0.34\text{ V}

) has the higher reduction potential, so it is correctly the cathode (reduction). Subtracting:

(+0.34) - (-0.76) = +1.10\text{ V}

. Positive result confirms spontaneity. Atom and charge balance check: left side has 1 Zn atom and 1 Cu²⁺ ion (net charge +2); right side has 1 Zn²⁺ ion (net charge +2) and 1 Cu atom. Atoms and charges are balanced. ✓

Common mistakes and how to avoid them

Flipping the sign of the anode's reduction potential manually before subtracting, then subtracting again. This double-corrects the sign and gives the wrong answer.

Correction: Look up both half-reactions as reduction potentials from the Data Booklet. Apply the formula directly:

E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}

. The subtraction already accounts for the reversal of the anode half-reaction. Do not flip any sign before substituting.

Multiplying the E° value by the same integer used to balance electrons in the half-reaction, treating E° like an enthalpy value that scales with moles.

Correction: Standard reduction potentials do not scale with the number of moles. If you multiply a half-reaction by 2 to balance electrons, the E° value stays exactly the same. Only the coefficients in the half-reaction equation change.

Assigning cathode and anode based on which species appears first in the equation or which has the larger absolute value of E°, rather than which has the more positive E° value.

Correction: Always assign the cathode to the half-reaction with the numerically higher (more positive) E° value. Compare the signed values: for example,

-0.28\text{ V}

is higher than

-0.76\text{ V}

, so the half-reaction at

-0.28\text{ V}

is the cathode.

Concluding that a reaction with a large negative

E°_{\text{cell}}

is spontaneous because the number looks large.

Correction: Only a positive

E°_{\text{cell}}

indicates a spontaneous reaction. A negative value — no matter how large in absolute terms — means the forward reaction is non-spontaneous. The reverse reaction would be spontaneous instead.

Using

E°_{\text{cell}} = E°_{\text{anode}} - E°_{\text{cathode}}

(reversing the subtraction order).

Correction: The correct Alberta Chemistry 30 Data Booklet formula is

E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}

. Cathode comes first in the subtraction. Reversing gives the opposite sign, which reverses the spontaneity conclusion.

Lesson summary

Standard reduction potentials (E°, in volts) measure how strongly a species attracts electrons when written as a reduction half-reaction. All values in the Alberta Chemistry 30 Data Booklet are measured relative to the standard hydrogen electrode ( $E° = 0.00\text{ V}$ ).
The half-reaction with the higher (more positive) E° is always assigned to the cathode (reduction), and the half-reaction with the lower E° is reversed and assigned to the anode (oxidation).
The standard cell potential is calculated as $E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}$ , using reduction potential values directly from the Data Booklet without changing their signs beforehand.
A positive $E°_{\text{cell}}$ means the forward reaction is spontaneous under standard conditions; a negative $E°_{\text{cell}}$ means it is non-spontaneous as written.
The E° value of a half-reaction does not change when the half-reaction is multiplied by a whole number to balance electron transfer.
At the particle level, spontaneous electron transfer flows from the species with the lower reduction potential (stronger reducing agent) to the species with the higher reduction potential (stronger oxidizing agent).

Check your understanding

Question 1

Using the Data Booklet values

E°(\text{Cu}^{2+}/\text{Cu}) = +0.34\text{ V}

and

E°(\text{Ag}^+/\text{Ag}) = +0.80\text{ V}

, what is

E°_{\text{cell}}

for the reaction

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

, and is it spontaneous?

$E°_{\text{cell}} = +0.46\text{ V}$ ; spontaneous
$E°_{\text{cell}} = -0.46\text{ V}$ ; non-spontaneous
$E°_{\text{cell}} = +1.14\text{ V}$ ; spontaneous
$E°_{\text{cell}} = +0.46\text{ V}$ ; non-spontaneous

Show answer and explanation

E°_{\text{cell}} = +0.46\text{ V}

; spontaneous

Silver ions have the higher E°

(+0.80 V), so

\text{Ag}

^+ is reduced at the cathode. Copper is oxidized at the anode (

E° = +0.34\text{ V}

). Using the formula:

E°_{\text{cell}} = (+0.80\text{ V}) - (+0.34\text{ V}) = +0.46\text{ V}

. Because

E°_{\text{cell}} > 0

, the reaction is spontaneous. Note: the half-reaction for silver is multiplied by 2 to balance electrons, but the E° value stays at +0.80 V, not doubled.

Question 2

A student wants to know whether magnesium metal will spontaneously reduce zinc ions in solution. Data Booklet values:

E°(\text{Zn}^{2+}/\text{Zn}) = -0.76\text{ V}

and

E°(\text{Mg}^{2+}/\text{Mg}) = -2.37\text{ V}

. Which conclusion is correct?

The reaction is non-spontaneous because both E° values are negative.
The reaction is spontaneous; $E°_{\text{cell}} = +1.61\text{ V}$ .
The reaction is non-spontaneous; $E°_{\text{cell}} = -1.61\text{ V}$ .
The reaction is spontaneous; $E°_{\text{cell}} = -3.13\text{ V}$ .

Show answer and explanation

The reaction is spontaneous;

E°_{\text{cell}} = +1.61\text{ V}

.

Zinc ions have the higher reduction potential (−0.76 V > −2.37 V), so

\text{Zn}^{2+}

is reduced at the cathode. Magnesium is oxidized at the anode.

E°_{\text{cell}} = (-0.76\text{ V}) - (-2.37\text{ V}) = +1.61\text{ V}

. Positive result means spontaneous. The fact that both values are negative does not make the reaction non-spontaneous; only the sign of

E°_{\text{cell}}

matters.

Question 3

A half-reaction is multiplied by 3 to balance electron transfer in an overall redox equation. What happens to its standard reduction potential E°?

It is multiplied by 3.
It is divided by 3.
It stays the same.
It becomes negative.

Show answer and explanation

It stays the same.

Standard reduction potentials are intensive quantities — they do not depend on the amount of substance. Multiplying a half-reaction by any integer changes only the coefficients, not the E° value. Always use the E° value exactly as listed in the Data Booklet.

Question 4

Given the two half-reactions below, predict whether the overall reaction as written is spontaneous.

Overall:

\text{Fe}(s) + 2\,\text{H}^+(aq) \rightarrow \text{Fe}^{2+}(aq) + \text{H}_2(g)

Data Booklet:

E°(\text{Fe}^{2+}/\text{Fe}) = -0.44\text{ V}

;

E°(\text{H}^+/\text{H}_2) = 0.00\text{ V}

Non-spontaneous; $E°_{\text{cell}} = -0.44\text{ V}$ .
Spontaneous; $E°_{\text{cell}} = +0.44\text{ V}$ .
Spontaneous; $E°_{\text{cell}} = +0.44\text{ V}$ , and Fe is the oxidizing agent.
Non-spontaneous; $E°_{\text{cell}} = 0.00\text{ V}$ .

Show answer and explanation

Spontaneous;

E°_{\text{cell}} = +0.44\text{ V}

.

Hydrogen ions have the higher reduction potential (0.00 V > −0.44 V), so

\text{H}^+

is reduced at the cathode and iron is oxidized at the anode.

E°_{\text{cell}} = (0.00\text{ V}) - (-0.44\text{ V}) = +0.44\text{ V}

. Positive, so the reaction is spontaneous. Option C is wrong because Fe is the reducing agent (it is oxidized), not the oxidizing agent.

Key terms

Standard reduction potential (E°): A value in volts that measures the tendency of a species to gain electrons (be reduced) under standard conditions (1 mol/L, 101.3 kPa, 25°C), relative to the standard hydrogen electrode.
Standard hydrogen electrode (SHE): The reference half-reaction $2\,\text{H}^+(aq) + 2\,e^- \rightarrow \text{H}_2(g)$ assigned $E° = 0.00\text{ V}$ , against which all other reduction potentials are measured.
Standard cell potential ( $E°_{\text{cell}}$ ): The overall voltage of a redox reaction under standard conditions, calculated as $E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}$ . A positive value indicates a spontaneous reaction.
Cathode: The electrode at which reduction (gain of electrons) occurs. In spontaneity calculations, the cathode is the half-reaction with the higher (more positive) standard reduction potential.
Anode: The electrode at which oxidation (loss of electrons) occurs. In spontaneity calculations, the anode is the half-reaction with the lower (more negative) standard reduction potential.
Spontaneous reaction: A reaction that proceeds on its own in the forward direction without a continuous input of energy. Under standard conditions, a redox reaction is spontaneous when $E°_{\text{cell}} > 0$ .
Oxidizing agent: The species that gains electrons in a redox reaction; it causes another species to be oxidized. The oxidizing agent is itself reduced, and it has the higher standard reduction potential.
Reducing agent: The species that loses electrons in a redox reaction; it causes another species to be reduced. The reducing agent is itself oxidized, and it has the lower standard reduction potential.

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30-B1.6k · Predict redox spontaneity using standard reduction potentials

Alberta Grade 12 Chemistry

Alberta Chemistry 30 — Outcome 30-B1.6k

What you will learn

Prerequisite Review: Oxidation, Reduction, and Half-Reactions

Standard Reduction Potentials: What They Mean

Calculating the Standard Cell Potential

Connecting the Particle Model to the Numbers

Applying the Method: Step-by-Step Strategy

Standard Reduction Potentials for Selected Half-Reactions (from Alberta Chemistry 30 Data Booklet)

Predicting Spontaneity for the Zinc–Copper Reaction

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