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F. Electrochemistry · Electrochemistry

Electrolytic Cells and Electrolysis

F2 F2.1 F2.6 F3.5About 75 minutes
Learning goals
  • I can explain how an electrolytic cell uses electrical energy to drive a non-spontaneous redox reaction.
  • I can compare galvanic and electrolytic cells, including the signs of the electrodes.
  • I can predict the products of electrolysis of molten salts and aqueous solutions.
  • I can calculate the minimum voltage needed for an electrolysis.

Words to know

electrolysis
Using electrical energy from an outside source to drive a non-spontaneous redox reaction.
electrolytic cell
A cell in which an external power supply forces a non-spontaneous reaction to occur. The anode is positive and the cathode is negative.
minimum (decomposition) voltage
The smallest voltage, in theory, that must be applied to make an electrolysis happen: equal to the size of the (negative) E°cell.
overvoltage (overpotential)
The extra voltage, beyond the theoretical value, that some half-reactions (especially those making gases such as O₂ and H₂) need in practice.
molten
Melted. In a molten ionic compound the ions can move, so it conducts electricity.

Running a reaction backwards

In the lesson Calculating Cell Potentials and Predicting Spontaneity, a negative E°cell meant "this reaction will not happen by itself". But "will not happen by itself" is not the same as "impossible". If you push electrons hard enough with a power supply, you can force a non-spontaneous reaction to go. That is electrolysis, and the setup is an electrolytic cell.

Electrolysis is how we make aluminium, sodium, chlorine, very pure copper and most "green" hydrogen. Every time you charge a phone, its battery works as an electrolytic cell.

Remember from Grade 11: in SCH3U (Chemical Reactions) you wrote decomposition reactions driven by electricity, such as 2NaCl(l) → 2Na(l) + Cl₂(g) and 2H₂O(l) → 2H₂(g) + O₂(g). Now you will see what happens at each electrode, and why.

molten NaCl(l), above 801 °CDC power−+e⁻e⁻cathode (−)reductionanode (+)oxidationNa⁺Cl⁻Na⁺ + e⁻ → Na(l)liquid sodium2Cl⁻ → Cl₂(g)+ 2e⁻chlorine gas
Electrolysis of molten sodium chloride with inert graphite electrodes. The power supply pumps electrons onto the cathode, where Na⁺ ions are reduced to sodium metal, and pulls electrons from the anode, where Cl⁻ ions are oxidized to chlorine gas.

How an electrolytic cell works

  • The power supply acts like an electron pump. Its negative terminal pushes electrons onto one electrode; its positive terminal pulls electrons off the other.
  • The electrode connected to the negative terminal has extra electrons, so reduction happens there: it is the cathode.
  • The electrode connected to the positive terminal loses electrons, so oxidation happens there: it is the anode.
  • Cations move toward the cathode and anions move toward the anode, just as in a galvanic cell.

The definitions never change (An Ox, Red Cat), but the signs flip: in an electrolytic cell the anode is positive and the cathode is negative.

GALVANIC CELLELECTROLYTIC CELLspontaneous reactionnon-spontaneous reactionE°cell > 0E°cell < 0 (needs power)chemical → electrical energyelectrical → chemical energyanode: oxidation, NEGATIVEanode: oxidation, POSITIVEcathode: reduction, POSITIVEcathode: reduction, NEGATIVEusually two compartmentsoften one containere.g. a battery in usee.g. charging a battery
Galvanic and electrolytic cells side by side. Oxidation is always at the anode and reduction at the cathode; what changes is whether the reaction is spontaneous and which electrode is positive.
Worked example

Example 1: Molten sodium chloride

Write the half-reactions and the overall reaction for the electrolysis of molten NaCl, and calculate the minimum voltage needed.

  1. Molten NaCl contains only Na⁺ and Cl⁻ ions: no water.
  2. Cathode (reduction): Na⁺ + e⁻ → Na, E° = −2.71 V.
  3. Anode (oxidation): 2Cl⁻ → Cl₂ + 2e⁻; as a reduction, E° = +1.36 V.
  4. Overall (sodium ×2): 2Na⁺ + 2Cl⁻ → 2Na + Cl₂.
  5. E°cell = E°cathode − E°anode = −2.71 − 1.36 = −4.07 V. Negative, so not spontaneous: a power supply must provide at least 4.07 V.

2NaCl(l) → 2Na(l) + Cl₂(g), minimum about 4.07 V. (These E° values are for aqueous solutions at 25 °C, so for a molten salt at 801 °C they only give an estimate; industrial cells run at about 7 V because of resistance and other losses.)

Electrolysis of aqueous solutions: water competes

In a solution, water can also be reduced or oxidized. At each electrode, the reaction that happens is the easiest one:

  • At the cathode, the strongest oxidizing agent present is reduced (highest E°).
  • At the anode, the strongest reducing agent present is oxidized (lowest E° as written in the table).

For water in a neutral solution (pH 7), use these values instead of the 1.0 mol/L values in the table:

  • Water reduced: 2H₂O + 2e⁻ → H₂ + 2OH⁻, E = −0.41 V at pH 7
  • Water oxidized: 2H₂O → O₂ + 4H⁺ + 4e⁻, E = +0.82 V (as a reduction) at pH 7

So, roughly: metal ions above about −0.41 V (Cu²⁺, Ag⁺) are reduced at the cathode; below it (Na⁺, K⁺, Mg²⁺, Al³⁺), water is reduced instead and hydrogen bubbles off. Anions with E° below +0.82 V (such as I⁻, +0.54 V) are oxidized at the anode instead of water. Sulfate and nitrate are not oxidized at all, so oxygen forms. Bromide and chloride are a special case (see the tip below).

Overvoltage: when the table is not enough

Making a gas at an electrode usually needs extra voltage, called overvoltage. Oxygen has a large overvoltage. That is why electrolysis of concentrated salt water (brine) gives chlorine, not oxygen, at the anode (concentrated bromide solutions likewise give bromine), even though the numbers (+1.36 V versus +0.82 V) say water should win. Hydrogen also has an overvoltage on some metals, which is why zinc can be electroplated from water even though Zn²⁺ is slightly below water in the table. On tests, apply the simple rule and the chloride exception unless told otherwise.

Worked example

Example 2: Aqueous potassium iodide

Predict the products of electrolysis of 1.0 mol/L KI(aq) with inert electrodes, write the overall reaction and find the minimum voltage (use the pH 7 values for water).

  1. Species: K⁺, I⁻, H₂O.
  2. Cathode: compare K⁺ (−2.93 V) and H₂O (−0.41 V). Water is the stronger oxidizing agent: 2H₂O + 2e⁻ → H₂ + 2OH⁻.
  3. Anode: compare I⁻ (+0.54 V) and H₂O (+0.82 V). Iodide is the stronger reducing agent: 2I⁻ → I₂ + 2e⁻.
  4. Overall: 2I⁻(aq) + 2H₂O(l) → I₂(aq) + H₂(g) + 2OH⁻(aq).
  5. E°cell = −0.41 − 0.54 = −0.95 V, so the minimum voltage is about 0.95 V.

Hydrogen gas and hydroxide ions at the cathode (phenolphthalein would turn pink there), brown iodine at the anode. Minimum about 0.95 V.

Worked example

Example 3: Copper(II) sulfate with inert electrodes

Predict the products when CuSO₄(aq) is electrolyzed with platinum electrodes.

  1. Cathode: Cu²⁺ (+0.34 V) beats water (−0.41 V): Cu²⁺ + 2e⁻ → Cu. Copper plates onto the cathode.
  2. Anode: SO₄²⁻ cannot be oxidized further (sulfur is already +6), so water is oxidized: 2H₂O → O₂ + 4H⁺ + 4e⁻.
  3. Overall (copper ×2): 2Cu²⁺ + 2H₂O → 2Cu + O₂ + 4H⁺.

Copper at the cathode, oxygen at the anode, and the solution becomes more acidic. The blue colour fades as Cu²⁺ is used up.

Worked example

Example 4: Splitting water

Pure water barely conducts, so a little sodium sulfate is added. Predict the products and the theoretical minimum voltage.

  1. Cathode: Na⁺ (−2.71 V) versus water (−0.41 V): water is reduced to H₂.
  2. Anode: SO₄²⁻ is not oxidized; water is oxidized to O₂.
  3. Overall: 2H₂O(l) → 2H₂(g) + O₂(g). The sodium sulfate only carries the current.
  4. E°cell = −0.41 − 0.82 = −1.23 V.

Hydrogen and oxygen, in a 2 : 1 volume ratio. At least 1.23 V in theory; practical electrolyzers need roughly 1.8 to 2 V because of overvoltage and resistance.

Just for fun

Sodium was first isolated in 1807 by Humphry Davy, using electrolysis of molten sodium hydroxide. He reportedly danced around his lab with joy. Very few chemists have danced since, but electrolysis is a good reason to try.

Hands-on activity

Kitchen electrolysis with baking soda

A grown-up must supervise. You need: a 9 V battery, 2 pencils sharpened at both ends (the graphite is your inert electrode), 2 alligator-clip leads (or insulated wire and tape), a clear glass, warm water, 1 teaspoon of baking soda (sodium hydrogen carbonate), and a piece of cardboard to hold the pencils.

Use baking soda only. Do not use table salt: salt water makes chlorine gas at the anode.

  1. Dissolve the baking soda in the glass of warm water.
  2. Push the pencils through the cardboard about 3 cm apart and rest the card on the glass so the lower pencil tips are in the water and not touching.
  3. Clip one lead from each battery terminal to the top graphite tip of a pencil.
  4. Watch the lower tips for 2 to 3 minutes. Which tip makes more bubbles?
  5. Disconnect the battery when finished. Pour the solution down the sink.

What to look for: bubbles form at both tips, with more at the tip connected to the negative terminal. That is the cathode, where hydrogen forms (2H₂O + 2e⁻ → H₂ + 2OH⁻). The positive tip is the anode, making oxygen, about half the volume. The gas amounts are tiny, but keep flames away anyway. Do not leave the battery connected for long: it warms up and runs down.

Your turn

1

In an electrolytic cell, which electrode is positive?

2

Write the half-reactions for the electrolysis of molten magnesium chloride.

3

Calculate the minimum voltage for electrolysis of molten MgCl₂ using table values.

4

Predict the products of electrolysis of AgNO₃(aq) with inert electrodes.

5

Predict the cathode product of electrolysis of NaNO₃(aq).

6

Why can't sodium metal be made by electrolysing NaCl(aq)?

7

Why is chlorine, not oxygen, produced at the anode when concentrated brine is electrolysed?

More practice: warm-up

1

Is electrolysis spontaneous?

2

At which electrode does reduction occur in an electrolytic cell?

3

Toward which electrode do anions move in electrolysis?

4

What gas forms at the cathode when water is electrolyzed?

5

What is the volume ratio of H₂ to O₂ when water is electrolyzed?

More practice: on your own

1

For the electrolysis of concentrated ZnBr₂(aq), first write what the simple E° rule predicts at each electrode, then what actually happens.

2

Calculate the minimum voltage to electrolyze CuCl₂(aq) to copper and chlorine.

3

A rechargeable battery is being charged. Is it acting as a galvanic or electrolytic cell? Which electrode is positive?

4

In the electrolysis of KI(aq), a few drops of phenolphthalein are added near each electrode. Where does the solution turn pink, and why?

5

Explain the difference between the theoretical minimum voltage and the voltage actually used in industry.

Challenge

1

A 1.0 mol/L solution contains Cu²⁺, Ag⁺ and Zn²⁺ ions. As the applied voltage is slowly increased, in what order do the metals plate onto the cathode? Explain.

2

Explain why electrolysis of NaCl(aq) produces a basic solution while electrolysis of CuSO₄(aq) produces an acidic one.

3

In the KI(aq) electrolysis from Example 2, a student applies 0.80 V and sees nothing happen. A second student applies 1.5 V and sees bubbles and brown colour. Explain both observations.

Check your understanding

Quiz

1.

In an electrolytic cell, the cathode is:

2.

What is produced at the cathode during electrolysis of molten KBr?

3.

What is produced at the cathode during electrolysis of aqueous KBr?

4.

Which ion is NOT oxidized at the anode during electrolysis of its aqueous salt, so that oxygen forms instead?

5.

The theoretical minimum voltage to electrolyze molten ZnCl₂ (E° Zn²⁺/Zn = −0.76 V, Cl₂/Cl⁻ = +1.36 V) is:

6.

Why does electrolysis of concentrated brine produce chlorine rather than oxygen at the anode?

For parents and teachers

Big idea (F2, F3.5): an electrolytic cell uses an outside power supply to force a non-spontaneous redox reaction. Oxidation is still at the anode and reduction at the cathode, but the anode is now positive. In solutions, water competes with the ions, and the easiest reduction and easiest oxidation win.

How to help: Do the baking-soda activity together, and ask your teen to predict which pencil will bubble more and why. Ask them to explain why salt must not be used.

Common difficulties:

  • Getting the electrode signs backwards. Remind them: the definitions (anode = oxidation) never change; only the signs do.
  • Forgetting that water competes in aqueous electrolysis.
  • Forgetting the overvoltage exception for concentrated chloride.
  • Writing a positive minimum voltage as "E°cell". E°cell is negative for electrolysis; the minimum applied voltage is its size.

Safety: use baking soda only (never salt, which makes chlorine). Supervise, keep flames away from the tiny amounts of hydrogen, and do not leave the 9 V battery connected for long.