F2.5 Draw and analyse labelled galvanic-cell diagrams | Ontario Grade 12 Chemistry (SCH4U)

Ontario Grade 12 Chemistry

Electrochemistry

SCH4U study topic F2.5: follow electron flow and identify each part of the cell

A galvanic cell can make a connected device respond, such as a small light turning on. The cell changes chemical energy into electrical energy through a reaction that transfers electrons. The reaction is separated into two locations so electrons travel through a wire rather than directly from one reacting particle to another. To read or draw a cell diagram, track where electrons are released, where they are accepted, and how ions move through the solutions and salt bridge.

What you will learn

Describe how a galvanic cell produces an observable electrical effect.
Identify the anode, cathode, electrodes, solutions, salt bridge, and electron-flow direction in a labelled diagram.
Use particle changes and half-reactions to analyse what happens at each electrode.
Draw a clear galvanic-cell diagram and interpret its cell notation.

1. From particle changes to a cell

A redox reaction transfers electrons. Oxidation means losing electrons; reduction means gaining electrons. These definitions are useful because they let you track electrons even when no visible electron is shown in a chemical equation.

In a galvanic cell, oxidation and reduction occur in separate half-cells. A half-cell is one electrode in contact with a solution containing ions that can react at that electrode. The electrode is a conducting solid, often a metal. The two electrodes are joined by a wire, and the two solutions are connected by a salt bridge: a material containing mobile ions that completes the circuit.

At the particle level, oxidation releases electrons at one electrode. Those electrons travel through the external wire to the other electrode, where particles in solution accept them during reduction. Ions move through the salt bridge to prevent charge from building up in either half-cell. The salt bridge does not carry electrons through the solutions.

The cell diagram should make these connections clear. Label both electrodes, both solutions, the salt bridge, the external wire, and the direction of electron flow. Include the anode and cathode names, not just the names of the metals.

\text{oxidation at anode} \rightarrow \text{reduction at cathode}

Oxidation releases electrons; reduction accepts electrons.
The anode is the electrode where oxidation occurs.
The cathode is the electrode where reduction occurs.
Electrons travel in the wire from anode to cathode.

2. Labels and charge movement

The anode and cathode are named by the reactions that occur there. In a galvanic cell, the anode releases electrons into the wire, so it is the negative electrode. The cathode receives electrons from the wire, so it is the positive electrode. These signs describe the galvanic cell while it is operating.

A common example pairs zinc metal with a solution containing copper(II) ions. Zinc atoms become zinc ions and release electrons at the zinc electrode. Copper(II) ions gain those electrons and form copper atoms at the copper electrode. The zinc electrode is therefore the anode, and the copper electrode is the cathode.

The salt bridge allows ions to move between the half-cells and helps maintain electrical neutrality. As zinc ions enter the anode solution, that solution gains positive charge. Negative ions from the bridge can move toward it. As copper(II) ions are removed from the cathode solution, positive ions from the bridge can move toward that side. The exact ions depend on the salt used.

The bridge completes the circuit by allowing ion movement. It does not provide the electrons for the electrode reactions. Electrons travel through the wire. In a diagram, show these as separate paths and label the electron-flow arrow from anode to cathode.

\mathrm{Zn(s)} \rightarrow \mathrm{Zn^{2+}(aq)} + 2e^-;\quad \mathrm{Cu^{2+}(aq)} + 2e^- \rightarrow \mathrm{Cu(s)}

For a galvanic cell, the anode is negative and the cathode is positive.
The wire carries electrons; the salt bridge permits ion movement.
A metal electrode may gain or lose solid metal as its atoms or ions react.

3. Representing a cell

A labelled drawing is often the clearest way to show physical parts and electron flow. Draw two separate beakers or half-cells. Put one electrode in each solution. Connect the electrodes with a wire and the solutions with a salt bridge. Mark the anode and cathode, the electrode signs, and the electron direction. Add the names and states of the electrode materials and solution ions.

Cell notation is a compact symbolic way to record the same arrangement. A single vertical line separates two phases in contact, such as a solid electrode and its aqueous ions. A double vertical line represents the salt bridge. By convention, the anode is written on the left and the cathode on the right.

For the zinc–copper cell, the notation starts with solid zinc, followed by aqueous zinc ions, then the salt bridge, then aqueous copper(II) ions and solid copper. This arrangement shows which substances are in each half-cell. It does not replace a labelled drawing when you need to show the wire, electrode signs, or electron-flow arrow.

Use the notation and the particle changes together. The left-hand half-cell must match oxidation, and the right-hand half-cell must match reduction. Check that electrons lost in the oxidation half-reaction equal electrons gained in the reduction half-reaction before adding the half-reactions.

\mathrm{Zn(s)}\,|\,\mathrm{Zn^{2+}(aq)}\,\|\,\mathrm{Cu^{2+}(aq)}\,|\,\mathrm{Cu(s)}

One vertical line marks a boundary between phases.
A double vertical line represents the salt bridge.
Write the anode half-cell on the left and the cathode half-cell on the right.
A complete cell drawing also labels the wire and electron flow.

4. A reliable analysis routine

Start by identifying the substances at each electrode and in each solution. Ask which particles lose electrons and which gain them. A metal electrode that forms aqueous positive ions is being oxidized. A positive ion that forms a neutral metal atom is being reduced.

Next, label anode and cathode from the reactions, not from left-versus-right position in an unfamiliar drawing. Then add electrode signs appropriate for a galvanic cell. Trace the electron arrow through the wire from the oxidation site to the reduction site.

Finally, inspect the salt bridge and the overall reaction. The bridge joins the solutions and allows ions to move. The overall reaction must conserve atoms and net charge. If the two half-reactions show different numbers of electrons, multiply a half-reaction as needed so that the electrons cancel when the reactions are added.

Use the reactions to identify electrodes; do not guess from their position.
Check electron transfer and conservation of atoms and charge.
Do not draw electrons travelling through the salt bridge.

Worked example

Analyse and label a zinc–copper cell

A cell contains a zinc strip in aqueous zinc nitrate and a copper strip in aqueous copper(II) sulfate. The half-cells are joined by a wire and a salt bridge. Identify the anode and cathode, electrode signs, electron-flow direction, half-reactions, and cell notation.

Identify the particle changes
Zinc atoms can enter solution as zinc ions, releasing electrons. Copper(II) ions can accept electrons and form copper metal. Therefore zinc is oxidized and copper(II) is reduced.
$\mathrm{Zn(s)} \rightarrow \mathrm{Zn^{2+}(aq)} + 2e^-;\quad \mathrm{Cu^{2+}(aq)} + 2e^- \rightarrow \mathrm{Cu(s)}$
Name and sign the electrodes
Oxidation occurs at the anode, so the zinc strip is the anode and is negative in this galvanic cell. Reduction occurs at the cathode, so the copper strip is the cathode and is positive.
$\text{anode: Zn (−)};\quad \text{cathode: Cu (+)}$
Trace electrons and combine the reactions
The released electrons travel through the wire from zinc to copper. The half-reactions already involve two electrons each, so they cancel when added. The resulting reaction conserves both atoms and net charge.
$\mathrm{Zn(s)} + \mathrm{Cu^{2+}(aq)} \rightarrow \mathrm{Zn^{2+}(aq)} + \mathrm{Cu(s)}$
Write the cell notation
Put the anode on the left and the cathode on the right. Use one vertical line for each electrode–solution boundary and a double vertical line for the salt bridge.
$\mathrm{Zn(s)}\,|\,\mathrm{Zn^{2+}(aq)}\,\|\,\mathrm{Cu^{2+}(aq)}\,|\,\mathrm{Cu(s)}$

Answer: The zinc strip is the negative anode; the copper strip is the positive cathode. Electrons flow through the wire from zinc to copper. Label the salt bridge between the solutions and show ion movement through it, not electron flow.

Check: The net ionic equation has one zinc atom and one copper atom on each side. Its total charge is +2 on each side.

Common mistakes and how to avoid them

Labelling the anode as the site of reduction.

Correction: The anode is where oxidation occurs. The cathode is where reduction occurs.

Drawing electrons through the salt bridge.

Correction: Draw electron flow through the external wire. The salt bridge permits ions to move between solutions.

Assuming the left electrode is always the anode.

Correction: In conventional cell notation the anode is written on the left, but in a drawing identify each electrode from its reaction.

Adding half-reactions without checking the electrons.

Correction: Make the number of electrons lost equal the number gained, then cancel them and check atoms and net charge.

Lesson summary

A galvanic cell separates oxidation and reduction so electrons can travel through a wire.
Oxidation occurs at the anode; reduction occurs at the cathode.
In a galvanic cell, the anode is negative and the cathode is positive.
The salt bridge permits ion movement; it is not the electron path.
Cell notation places the anode on the left, the salt bridge in the middle, and the cathode on the right.

Check your understanding

Question 1

In a galvanic cell, where do electrons travel through the external wire?

From the cathode to the anode
From the anode to the cathode
From the salt bridge to the anode
From the solution directly into the salt bridge as electrons

Show answer and explanation

From the anode to the cathode

Oxidation at the anode releases electrons. Reduction at the cathode uses them.

Question 2

A diagram shows metal atoms becoming aqueous positive ions at one electrode. What should that electrode be labelled?

Cathode, because ions are formed
Anode, because oxidation occurs
Cathode, because electrons leave it
Salt bridge, because charge changes

Show answer and explanation

Anode, because oxidation occurs

The atoms lose electrons as they become positive ions. That is oxidation, which occurs at the anode.

Key terms

Anode: The electrode where oxidation occurs.
Cathode: The electrode where reduction occurs.
Electrode: A conducting solid where a half-reaction takes place.
Half-cell: An electrode together with the solution containing the reacting particles.
Salt bridge: A connection containing mobile ions that completes the circuit between half-cell solutions.
Cell notation: A compact symbolic representation of the two half-cells and the salt bridge.

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Published by DoAssignment. This reviewed lesson follows Ontario Grade 12 Chemistry (SCH4U), expectation F2.5. 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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F2.5 · Draw and analyse labelled galvanic-cell diagrams