Explain the standard hydrogen electrode reference potential

Alberta Grade 12 Chemistry

Electrochemical Changes

Alberta Chemistry 30 – Unit B2: Electrochemistry (Outcome 30-B2.5k)

Imagine connecting a voltmeter to a single metal strip dipped in a solution. The voltmeter reads nothing useful — a voltage is always a difference between two points, not a property of one point alone. Every measurement of an electrode's tendency to gain or lose electrons faces this same problem. Chemists solved it by agreeing on one universal reference point: the standard hydrogen electrode (SHE), whose potential is defined as exactly zero volts. All other electrode potentials are then measured and reported relative to this agreed-upon zero. This lesson explains what the SHE is, why its potential is set to zero, and how that choice lets chemists compare and use electrode potentials found in the Data Booklet.

What you will learn

Explain why a reference electrode is needed to measure electrode potentials.
Describe the construction and conditions of the standard hydrogen electrode (SHE).
State and apply the assigned potential of the SHE and explain why it is set to exactly zero.
Use the SHE as a reference to interpret standard reduction potential values from the Alberta Chemistry 30 Data Booklet.
Identify the half-reactions at the SHE when it acts as an anode or a cathode.

Prerequisite Review: Half-Reactions and Reduction Potentials

In Chemistry 20 you learned that oxidation is the loss of electrons and reduction is the gain of electrons. In an electrochemical cell, these two processes are separated into half-cells. Each half-cell contains a half-reaction — either an oxidation (at the anode) or a reduction (at the cathode).

A reduction potential measures how strongly a species pulls electrons toward itself under a given set of conditions. A large, positive reduction potential means the species is a strong oxidizing agent — it strongly attracts electrons. A large, negative reduction potential means the species strongly resists gaining electrons and prefers to give them up instead.

Standard conditions in Alberta Chemistry 30 mean: all dissolved ions at a concentration of

1\ \text{mol/L}

, all gases at a pressure of

101.325\ \text{kPa}

, and a temperature of

25\ ^\circ\text{C}

(298 K). The symbol

E^\circ

(read 'E-standard' or 'E-naught') means the reduction potential measured under these exact standard conditions. You will see

E^\circ

values in your Alberta Chemistry 30 Data Booklet reduction potential table.

\text{Oxidized form} + n\,e^- \rightarrow \text{Reduced form} \quad E^\circ

Oxidation = loss of electrons (occurs at the anode).
Reduction = gain of electrons (occurs at the cathode).
Standard conditions: $c = 1\ \text{mol/L}$ , $P = 101.325\ \text{kPa}$ , $T = 25\ ^\circ\text{C}$ .
$E^\circ$ is the standard reduction potential — always written as a reduction half-reaction.

The Problem: You Cannot Measure One Electrode Alone

Picture a zinc strip sitting in a

1\ \text{mol/L}

zinc sulfate solution. Zinc atoms at the surface have some tendency to release electrons and dissolve as

\text{Zn}^{2+}

ions. But how do you measure that tendency? A voltmeter needs two terminals. If you connect one terminal to the zinc strip, you still need something to connect the other terminal to — and whatever you choose will have its own tendency to gain or lose electrons. Your reading will always reflect both electrodes at once.

This is not a flaw in the equipment. It is a fundamental fact: electrical potential, like height above sea level, is always relative. Just as every elevation on Earth is measured relative to an agreed sea-level zero, every electrode potential must be measured relative to an agreed electrical zero.

Chemists around the world agreed to use a specific, reproducible electrode as that zero point. They call it the standard hydrogen electrode. By definition, its potential is set to exactly

0\ \text{V}

under standard conditions. All other

E^\circ

values in the Data Booklet are measured against this reference.

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

A single electrode potential cannot be measured in isolation — voltage is always a difference.
A universal reference electrode is needed so that measurements made in different labs are comparable.
The SHE is the internationally agreed reference with an assigned potential of exactly $0\ \text{V}$ .
All Data Booklet $E^\circ$ values are measured relative to the SHE.

Construction and Conditions of the Standard Hydrogen Electrode

The SHE is a carefully controlled half-cell. It consists of three components working together. First, a platinum electrode is used. Platinum is chosen because it does not react with hydrogen gas or with the acid solution; it simply provides a conducting surface where the electron-transfer reaction can occur. Second, hydrogen gas is bubbled over the platinum surface at a pressure of exactly

101.325\ \text{kPa}

(standard atmospheric pressure). Third, the platinum is immersed in an acidic solution containing hydrogen ions at a concentration of exactly

1\ \text{mol/L}

, typically prepared from hydrochloric acid or sulfuric acid.

Under these precise standard conditions, the hydrogen half-reaction reaches a dynamic equilibrium at the platinum surface. The half-reaction written as a reduction is:

\text{H}_2

molecules give up electrons to form

\text{H}^+

ions (oxidation direction) or

\text{H}^+

ions accept electrons to form

\text{H}_2

gas (reduction direction). Because the SHE can run in either direction depending on what it is paired with, it can act as either the anode or the cathode in a cell.

Think of the platinum as a reaction platform. Hydrogen gas molecules (

\text{H}_2

) adsorb onto its surface. There, they can either release electrons to become

\text{H}^+

ions (if the SHE is the anode), or

\text{H}^+

ions from the solution can accept electrons to form

\text{H}_2

gas (if the SHE is the cathode). The platinum itself does not appear in the net half-reaction equation because it is not consumed — it is a catalyst-like surface.

2\,\text{H}^+\text{(aq)} + 2\,e^- \rightarrow \text{H}_2\text{(g)} \quad E^\circ = 0\ \text{V}

SHE components: platinum electrode, $\text{H}_2$ gas at $101.325\ \text{kPa}$ , $[\text{H}^+] = 1\ \text{mol/L}$ .
Platinum is inert — it conducts electrons but does not react.
The SHE half-reaction: $2\,\text{H}^+\text{(aq)} + 2\,e^- \rightleftharpoons \text{H}_2\text{(g)}$ .
The SHE can act as either anode or cathode depending on its partner half-cell.
The assigned standard potential of the SHE is exactly $E^\circ = 0\ \text{V}$ .

Why Zero? Interpreting the Data Booklet Table

Assigning zero to the SHE is a human convention — like choosing sea level as the zero for altitude. It does not mean the SHE has no energy or no tendency to react. It simply means every other electrode is described by how it compares to the SHE. An electrode with

E^\circ = +0.80\ \text{V}

(such as

\text{Au}^{3+}/\text{Au}

) has a stronger pull on electrons than the SHE. An electrode with

E^\circ = -0.76\ \text{V}

(such as

\text{Zn}^{2+}/\text{Zn}

) has a weaker pull on electrons than the SHE — it prefers to give electrons to the SHE when the two are connected.

When a chemist pairs an unknown electrode with the SHE and measures the cell voltage, that measured voltage equals the unknown electrode's

E^\circ

value directly, because

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

and one of those terms is zero. If the unknown electrode is the cathode (reduction occurs there), the measured voltage is positive and equals

E^\circ_{\text{unknown}}

. This is exactly how the Data Booklet table was built.

The Data Booklet lists half-reactions in order from most negative

E^\circ

at the top (strongest reducing agents, weakest oxidizing agents) to most positive

E^\circ

at the bottom (strongest oxidizing agents, weakest reducing agents). The SHE entry sits in the middle of this table at

0\ \text{V}

, acting as the dividing line. Species above it in the table are better reducing agents than

\text{H}_2

; species below it are better oxidizing agents than

\text{H}^+

.

E°_{unknown} = E°_{cell} (when paired with SHE, which contributes 0\ V)

$E^\circ = 0\ \text{V}$ for the SHE is a defined, agreed-upon convention, not a measured value.
Positive $E^\circ$ : species gains electrons more readily than $\text{H}^+$ (stronger oxidizing agent than $\text{H}^+$ ).
Negative $E^\circ$ : species gains electrons less readily than $\text{H}^+$ (stronger reducing agent than $\text{H}_2$ ).
Pairing any electrode with the SHE lets you read its $E^\circ$ directly from the cell voltage.
The Data Booklet table is organized with the SHE at $0\ \text{V}$ as the reference midpoint.

Using the SHE to Predict Cell Behaviour

Knowing the SHE's role lets you use the Data Booklet confidently. When you connect any electrode to the SHE, the electrode with the more positive

E^\circ

will act as the cathode (site of reduction) and the electrode with the more negative

E^\circ

will act as the anode (site of oxidation). Because the SHE is at

0\ \text{V}

, any electrode with a positive

E^\circ

will act as the cathode when paired with the SHE, and any electrode with a negative

E^\circ

will act as the anode when paired with the SHE.

For example, a copper half-cell has

E^\circ_{\text{Cu}^{2+}/\text{Cu}} = +0.34\ \text{V}

. When paired with the SHE, copper is the cathode. Copper ions are reduced:

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

. At the SHE (now the anode), hydrogen gas is oxidized:

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

. The measured cell voltage is

+0.34\ \text{V} - 0\ \text{V} = +0.34\ \text{V}

, which is exactly the copper electrode's listed

E^\circ

value.

This confirms the fundamental purpose of the SHE: it is the universal measuring stick. Every

E^\circ

value in your Data Booklet was determined by pairing that half-cell with the SHE and measuring the resulting cell voltage.

E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = E^\circ_{\text{cathode}} - 0\ \text{V} = E^\circ_{\text{cathode}}

Electrode with more positive $E^\circ$ → cathode (reduction) when paired with SHE.
Electrode with more negative $E^\circ$ → anode (oxidation) when paired with SHE.
Any positive- $E^\circ$ electrode will be the cathode when paired with the SHE.
Any negative- $E^\circ$ electrode will be the anode when paired with the SHE.
The measured cell voltage equals the partner electrode's $E^\circ$ because the SHE contributes $0\ \text{V}$ .

Selected Standard Reduction Potentials Relative to the SHE (Alberta Chemistry 30 Data Booklet)

Half-Reaction (Reduction)	$E^\circ$ (V)	Interpretation Relative to SHE
$\text{Zn}^{2+}\text{(aq)} + 2\,e^- \rightarrow \text{Zn(s)}$	$-0.76$	Weaker pull on electrons than $\text{H}^+$ ; Zn is anode vs. SHE
$\text{Fe}^{2+}\text{(aq)} + 2\,e^- \rightarrow \text{Fe(s)}$	$-0.44$	Weaker pull on electrons than $\text{H}^+$ ; Fe is anode vs. SHE
$2\,\text{H}^+\text{(aq)} + 2\,e^- \rightarrow \text{H}_2\text{(g)}$	$0.00$	The SHE reference — defined as zero
$\text{Cu}^{2+}\text{(aq)} + 2\,e^- \rightarrow \text{Cu(s)}$	$+0.34$	Stronger pull on electrons than $\text{H}^+$ ; Cu is cathode vs. SHE
$\text{Ag}^+\text{(aq)} + e^- \rightarrow \text{Ag(s)}$	$+0.80$	Stronger pull on electrons than $\text{H}^+$ ; Ag is cathode vs. SHE

Worked example

Determining the Role of the SHE and Calculating Cell Voltage

A zinc half-cell operates under standard conditions: a zinc strip in

1\ \text{mol/L}

\text{ZnSO}_4\text{(aq)}

. It is connected to a standard hydrogen electrode. The standard reduction potential for zinc is

E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76\ \text{V}

(from the Alberta Chemistry 30 Data Booklet). (a) Identify which half-cell is the cathode and which is the anode. (b) Write the balanced half-reactions and the overall balanced cell reaction. (c) Calculate the standard cell potential.

Step 1 – Compare the two reduction potentials.
List both reduction potentials. The half-cell with the more positive (or less negative) reduction potential will act as the cathode. The half-cell with the more negative reduction potential will act as the anode.
$E^\circ_{\text{SHE}} = 0\ \text{V} \qquad E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76\ \text{V}$
Step 2 – Assign cathode and anode.
Since $0\ \text{V} > -0.76\ \text{V}$ , the SHE has the more positive reduction potential. Reduction (gain of electrons) occurs at the more positive electrode, so the SHE is the cathode. Oxidation occurs at the zinc electrode, so zinc is the anode.
$\text{Cathode: SHE} \quad (E^\circ = 0\ \text{V}) \qquad \text{Anode: Zn} \quad (E^\circ = -0.76\ \text{V})$
Step 3 – Write the cathode half-reaction (reduction at SHE).
At the SHE, hydrogen ions from solution gain electrons to form hydrogen gas. This is the reduction direction of the SHE half-reaction.
$2\,\text{H}^+\text{(aq)} + 2\,e^- \rightarrow \text{H}_2\text{(g)}$
Step 4 – Write the anode half-reaction (oxidation at zinc).
At the zinc anode, zinc metal loses electrons and dissolves as zinc ions. Reverse the reduction half-reaction listed in the Data Booklet to write it as oxidation.
$\text{Zn(s)} \rightarrow \text{Zn}^{2+}\text{(aq)} + 2\,e^-$
Step 5 – Check electron balance and write the overall cell reaction.
Both half-reactions involve exactly 2 electrons, so they balance directly. Add the two half-reactions and cancel the electrons on both sides.
$\text{Zn(s)} + 2\,\text{H}^+\text{(aq)} \rightarrow \text{Zn}^{2+}\text{(aq)} + \text{H}_2\text{(g)}$
Step 6 – Calculate the standard cell potential.
Use the standard cell potential formula. Subtract the anode's reduction potential from the cathode's reduction potential. Note that you use the reduction potential of the anode as listed — do not change its sign in the formula.
$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = 0\ \text{V} - (-0.76\ \text{V}) = +0.76\ \text{V}$

Answer: The SHE is the cathode and zinc is the anode. The overall balanced reaction is

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

. The standard cell potential is

E^\circ_{\text{cell}} = +0.76\ \text{V}

.

Check: A positive

E^\circ_{\text{cell}}

confirms a spontaneous reaction under standard conditions, which is consistent with zinc being a more active metal than hydrogen. Also,

+0.76\ \text{V}

is exactly the magnitude of the zinc reduction potential from the Data Booklet, which is expected: because the SHE contributes

0\ \text{V}

, the cell voltage equals the partner electrode's

E^\circ

value directly. Atoms and charges are balanced: 1 Zn on each side; 2 H on each side; net charge on left =

2+

, net charge on right =

2+

. ✓

Common mistakes and how to avoid them

Thinking

E^\circ = 0\ \text{V}

for the SHE means it has no tendency to react or no energy.

Correction: Zero volts is a chosen reference point, not an absence of reactivity. The SHE is chemically active — it simply serves as the agreed measuring stick against which all other potentials are compared, exactly as sea level (0 m) is not the absence of land.

Changing the sign of

E^\circ_{\text{anode}}

before substituting it into the formula

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

.

Correction: Always substitute the reduction potential values from the Data Booklet directly into the formula. The subtraction sign in the formula already accounts for the fact that the anode runs in reverse. For example:

E^\circ_{\text{cell}} = 0\ \text{V} - (-0.76\ \text{V}) = +0.76\ \text{V}

, not

0\ \text{V} - (+0.76\ \text{V})

.

Believing the platinum electrode in the SHE participates in the half-reaction and should appear in the chemical equation.

Correction: Platinum is inert. It provides a surface for electron transfer but is neither consumed nor produced. It does not appear in the half-reaction equation. Only

\text{H}^+

,

e^-

, and

\text{H}_2

appear.

Assuming the SHE is always the cathode no matter which electrode it is paired with.

Correction: The SHE can be either the cathode or the anode. If it is paired with an electrode that has a negative

E^\circ

(such as zinc), the SHE is the cathode. If it is paired with an electrode that has a positive

E^\circ

(such as copper), the SHE is the anode. Always compare reduction potentials to decide.

Forgetting to specify standard conditions when stating that the SHE has

E^\circ = 0\ \text{V}

.

Correction: The potential of 0 V applies only under standard conditions:

[\text{H}^+] = 1\ \text{mol/L}

,

P_{\text{H}_2} = 101.325\ \text{kPa}

, and

T = 25\ ^\circ\text{C}

. If these conditions change, the potential of the hydrogen electrode changes too — but that is beyond the scope of Chemistry 30.

Lesson summary

A single electrode potential cannot be measured in isolation because voltage is always a difference between two points.
The standard hydrogen electrode (SHE) is the internationally agreed reference electrode, assigned a potential of exactly $E^\circ = 0\ \text{V}$ by convention.
The SHE consists of a platinum electrode, hydrogen gas at $101.325\ \text{kPa}$ , and $[\text{H}^+] = 1\ \text{mol/L}$ at $25\ ^\circ\text{C}$ . Its half-reaction is $2\,\text{H}^+\text{(aq)} + 2\,e^- \rightarrow \text{H}_2\text{(g)}$ .
All standard reduction potentials ( $E^\circ$ values) in the Alberta Chemistry 30 Data Booklet are measured relative to the SHE. Positive values mean a stronger pull on electrons than $\text{H}^+$ ; negative values mean a weaker pull.
When paired with the SHE, an electrode with positive $E^\circ$ is the cathode and an electrode with negative $E^\circ$ is the anode. The measured cell voltage equals that electrode's $E^\circ$ value directly, since the SHE contributes $0\ \text{V}$ .

Check your understanding

Question 1

Why is the standard reduction potential of the SHE defined as exactly

0\ \text{V}

?

Because hydrogen gas has no electrons to donate or accept.
Because it is an internationally agreed reference point that allows all other electrode potentials to be compared on the same scale.
Because the platinum electrode cancels out any potential that hydrogen would otherwise have.
Because $\text{H}^+$ ions and $\text{H}_2$ gas have equal and opposite potentials that sum to zero.

Show answer and explanation

Because it is an internationally agreed reference point that allows all other electrode potentials to be compared on the same scale.

The value

0\ \text{V}

is a human convention — a chosen reference, just like sea level for altitude. It is not a physical property of hydrogen itself. Setting it to zero means every other electrode potential is automatically expressed as a difference from this agreed standard, making Data Booklet values universally comparable.

Question 2

Under standard conditions, a nickel half-cell (

E^\circ_{\text{Ni}^{2+}/\text{Ni}} = -0.25\ \text{V}

) is connected to a standard hydrogen electrode. What is the standard cell potential and which electrode is the anode?

$E^\circ_{\text{cell}} = -0.25\ \text{V}$ ; the SHE is the anode.
$E^\circ_{\text{cell}} = +0.25\ \text{V}$ ; the SHE is the cathode and nickel is the anode.
$E^\circ_{\text{cell}} = +0.25\ \text{V}$ ; nickel is the cathode and the SHE is the anode.
$E^\circ_{\text{cell}} = 0\ \text{V}$ ; neither electrode acts as an anode because the SHE is always the reference.

Show answer and explanation

E^\circ_{\text{cell}} = +0.25\ \text{V}

; the SHE is the cathode and nickel is the anode.

The SHE (

0\ \text{V}

) has the more positive reduction potential compared to nickel (

-0.25\ \text{V}

), so the SHE is the cathode and nickel is the anode. Using the formula:

E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = 0\ \text{V} - (-0.25\ \text{V}) = +0.25\ \text{V}

.

Question 3

Which of the following correctly describes the physical setup of the standard hydrogen electrode (SHE)?

A copper strip immersed in $1\ \text{mol/L}$ $\text{HCl(aq)}$ with oxygen gas bubbled over it at $101.325\ \text{kPa}$ .
A platinum strip immersed in $1\ \text{mol/L}$ $\text{H}^+\text{(aq)}$ with hydrogen gas bubbled over it at $101.325\ \text{kPa}$ .
A platinum strip immersed in pure water with hydrogen gas bubbled over it at $101.325\ \text{kPa}$ .
A zinc strip immersed in $1\ \text{mol/L}$ $\text{HCl(aq)}$ with hydrogen gas bubbled over it at $101.325\ \text{kPa}$ .

Show answer and explanation

A platinum strip immersed in

1\ \text{mol/L}

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

with hydrogen gas bubbled over it at

101.325\ \text{kPa}

.

The SHE requires all three standard components: an inert platinum electrode (not copper or zinc, which would react),

[\text{H}^+] = 1\ \text{mol/L}

(not pure water), and

\text{H}_2

gas at

101.325\ \text{kPa}

(not oxygen). Platinum is essential because it is inert and provides a surface for the half-reaction without participating in it.

Question 4

A silver half-cell (

E^\circ_{\text{Ag}^+/\text{Ag}} = +0.80\ \text{V}

) is paired with the SHE. Which statement correctly describes what happens at the SHE?

The SHE acts as the cathode; $\text{H}^+$ ions are reduced to $\text{H}_2$ gas.
The SHE acts as the cathode; $\text{H}_2$ gas is reduced to $\text{H}^-$ ions.
The SHE acts as the anode; $\text{H}_2$ gas is oxidized to $\text{H}^+$ ions.
The SHE acts as the anode; $\text{H}^+$ ions are oxidized to $\text{H}_2$ gas.

Show answer and explanation

The SHE acts as the anode;

\text{H}_2

gas is oxidized to

\text{H}^+

ions.

Silver (

+0.80\ \text{V}

) has a more positive reduction potential than the SHE (

0\ \text{V}

), so silver is the cathode and the SHE is the anode. At the anode, oxidation occurs:

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

. Hydrogen gas loses electrons — that is oxidation, not reduction — to become

\text{H}^+

ions.

Key terms

Standard hydrogen electrode (SHE): The internationally agreed reference half-cell consisting of a platinum electrode in contact with $\text{H}_2$ gas at $101.325\ \text{kPa}$ and $[\text{H}^+] = 1\ \text{mol/L}$ at $25\ ^\circ\text{C}$ . Its standard potential is defined as exactly $0\ \text{V}$ .
Reference electrode: An electrode whose potential is fixed and known, used as a comparison point so that the potential of any other electrode can be measured.
Standard reduction potential ( $E^\circ$ ): The voltage measured for a half-cell under standard conditions ( $1\ \text{mol/L}$ , $101.325\ \text{kPa}$ , $25\ ^\circ\text{C}$ ) when paired against the SHE. Always written for the reduction direction.
Cathode: The electrode where reduction (gain of electrons) occurs. In a galvanic cell, it is the electrode with the more positive reduction potential.
Anode: The electrode where oxidation (loss of electrons) occurs. In a galvanic cell, it is the electrode with the more negative reduction potential.
Standard conditions: The defined reference state for electrochemical measurements: all dissolved species at $1\ \text{mol/L}$ , all gases at $101.325\ \text{kPa}$ , and temperature at $25\ ^\circ\text{C}$ .
Inert electrode: An electrode, such as platinum, that conducts electricity and provides a surface for the half-reaction but does not itself participate in or get consumed by the reaction.
Convention: An agreed-upon rule or definition adopted by scientists worldwide to make measurements consistent and comparable. The SHE potential of $0\ \text{V}$ is a convention, not a measured value.

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Published by DoAssignment. This reviewed lesson follows Alberta Grade 12 Chemistry (CHEM30), expectation 30-B2.5k. It is a study resource, not an official curriculum publication.

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30-B2.5k · Explain the standard hydrogen electrode reference potential

Alberta Grade 12 Chemistry

Alberta Chemistry 30 – Unit B2: Electrochemistry (Outcome 30-B2.5k)

What you will learn

Prerequisite Review: Half-Reactions and Reduction Potentials

The Problem: You Cannot Measure One Electrode Alone

Construction and Conditions of the Standard Hydrogen Electrode

Why Zero? Interpreting the Data Booklet Table

Using the SHE to Predict Cell Behaviour

Selected Standard Reduction Potentials Relative to the SHE (Alberta Chemistry 30 Data Booklet)

Determining the Role of the SHE 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