How to construct an electrochemical cell diagram

A neat, correctly labelled electrochemical cell diagram is a reliable source of marks in Paper 2 and Paper 3, yet many students lose them by mixing up which terminal is positive or which way the electrons flow. This guide shows you how to build the diagram piece by piece and label every part the way an examiner expects, using the voltaic and electrochemical cells topic. Once the layout is a habit, you can draw it in under two minutes.

What a voltaic cell does

A voltaic (galvanic) cell turns a spontaneous redox reaction into electricity. Two different metals dip into electrolyte and are joined by a wire; because one metal loses electrons more readily than the other, electrons are pushed through the wire and a current flows. Building one by hand is a standard practical, see constructing a simple voltaic cell for the method.

The parts you must draw

Every full-mark cell diagram shows these labelled components:

• Two electrodes, the two different metals, each labelled with its name or symbol.
• The electrolyte(s), the solution each electrode dips into, labelled (for example copper(II) sulfate solution).
• A salt bridge, in a Daniell-type cell with two separate beakers, joining the two solutions.
• A voltmeter or galvanometer, in the external circuit, to read the voltage.
• Connecting wires with the terminals marked positive (+) and negative (−).
• An arrow showing electron flow in the external wire.

Leave any one of these off and you leave marks on the table.

Step-by-step: build the diagram

1. Draw the container(s). For a Daniell cell, two beakers side by side; for a simple cell, one beaker with both electrodes.
2. Add the two electrodes dipping in, one on each side, and write the metal on each.
3. Label each electrolyte under or beside its beaker.
4. Draw the salt bridge as an inverted U⁻tube linking the two solutions (Daniell type only).
5. Connect the electrodes with wires up to a voltmeter in the middle.
6. Decide the terminals using the electrochemical series (next section), and mark + and −.
7. Add the electron-flow arrow from the negative terminal, through the wire, to the positive terminal.

Which terminal is which

This is the step students get wrong, so slow down. Use the electrochemical series:

• The more electropositive (more reactive) metal is the negative terminal. It is oxidised, it loses electrons, so it is the anode.
• The less reactive metal is the positive terminal. It is the cathode, where reduction happens.
• Electrons flow in the external wire from the negative terminal to the positive terminal, that is, from the more reactive metal to the less reactive one.

A memory hook: the more reactive metal is keener to give electrons away, so it becomes negative and pushes electrons out into the wire.

Worked example: the Daniell cell

Take zinc and copper, the classic pair. Zinc sits above copper in the reactivity series, so:

• Zinc electrode, negative terminal (anode). It is oxidised:

Zn → Zn²⁺ + 2e⁻

• Copper electrode, positive terminal (cathode). Copper(II) ions are reduced:

Cu²⁺ + 2e⁻ → Cu

Electrons leave the zinc, travel through the wire and voltmeter to the copper, and reduce Cu²⁺ there. The zinc electrode slowly dissolves and gets thinner; the copper electrode gets a fresh copper coating and gets thicker. The salt bridge lets ions move to keep both solutions electrically neutral. This is exactly the direction predicted when the more reactive metal displaces the less reactive one in displacement of metals.

Reading and predicting the voltage

The further apart the two metals are in the electrochemical series, the larger the voltage the cell produces. A zinc–copper cell gives more voltage than a zinc–iron cell, because zinc and copper are further apart. You will not be asked to calculate an exact voltage in SPM, but you should be able to compare two cells and say which gives the higher reading, and why.

Labelling checklist before you move on

• Both electrodes named, both electrolytes named.
• Terminals marked + and −, matching the reactivity order.
• Electron-flow arrow from negative to positive in the external wire.
• Half equations written, with electrons on the correct side.
• Salt bridge shown if the solutions are separate.

Tick all five and the diagram is complete.

Common mistakes to avoid

• Putting the arrow inside the solution. Electron flow is drawn in the wire, not through the electrolyte, inside the cell, charge is carried by ions.
• Calling the negative electrode the cathode. In a voltaic cell the anode (oxidation) is negative; the cathode (reduction) is positive. This is the reverse of electrolysis, so keep the two topics separate in your head.
• Forgetting the voltmeter or salt bridge. Both are easy marks.

Practise the layout until it is automatic

Draw the Daniell cell from memory a few times, then try other metal pairs and predict the terminals and half equations for each. The apparatus never really changes, only the metals and the electrolytes. If the anode/cathode and positive/negative labels keep flipping in your mind, that is a common tangle we untangle quickly in lessons; our online one-to-one classes run in English from RM50 an hour, with a paid one-hour trial. See how it works if you would like a teacher to check your diagrams against real Paper 2 answers.