Chemistry (0620) | Electrolysis

Electrolysis: Harnessing Electricity to Make Chemical Reactions Happen

Hello future chemists! This chapter, Electrochemistry, connects two of the most important fields in science: electricity and chemical reactions. We are focusing specifically on Electrolysis, a powerful technique used globally, from refining metals to manufacturing crucial chemicals like chlorine and aluminium.

Don't worry if this seems tricky at first—electrolysis is just forced chemistry! We will break down the rules for predicting what happens at each end of the circuit.

4.1 The Core Concepts of Electrolysis

What is Electrolysis? (Core 4.1.1)

Definition: Electrolysis is the decomposition (breaking down) of an ionic compound, either when molten or in an aqueous solution, by the passage of an electric current.

• It is a non-spontaneous reaction, meaning it requires energy (electricity) to happen.
• It involves redox reactions: oxidation happens at one electrode, and reduction happens at the other.

The Electrolytic Cell Setup (Core 4.1.2)

An electrolytic cell requires three key components:

1. Electrolyte: The molten or aqueous ionic substance that is decomposed. It must be able to conduct electricity via the movement of ions. (Remember: Solids cannot be electrolytes because their ions are fixed in a lattice.)
2. Electrodes: Two rods (conductors) inserted into the electrolyte, connected to the power source.
3. Power Source: Provides the direct current (DC) needed to drive the reaction.

Quick Memory Aid: The AC/DC Rules

• Anode: Positive electrode (PA)
• Cathode: Negative electrode (NC)

Analogy: Think of the electrodes as charging stations. The positive electrode (anode) attracts the negative ions, and the negative electrode (cathode) attracts the positive ions.

Transfer of Charge (Supplement 4.1.8)

For electrolysis to occur, charge must flow:

1. In the External Circuit (Wires): Electrons (\(e^-\)) flow from the anode to the cathode (pushed by the power source).
2. In the Electrolyte (Liquid/Solution): The ions carry the charge.
   • Cations (positive ions) move towards the Cathode (negative electrode).
   • Anions (negative ions) move towards the Anode (positive electrode).
3. At the Electrodes: The chemical reactions happen.
   • At the Cathode, Reduction occurs (gain of electrons).
   • At the Anode, Oxidation occurs (loss of electrons).

Memory Aid: OIL RIG
Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)

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Section 2: Electrolysis of Molten Compounds

The electrolysis of a molten binary compound is the simplest case because there are only two types of ions present.

Predicting Products for Molten Compounds (Core 4.1.5)

When an ionic compound is melted, there are no water ions (\(H^+\) or \(OH^-\)) to compete with the salt ions. Therefore, the products are always the constituent elements:

1. At the Cathode (Negative): The metal cation is reduced to the pure metal.
2. At the Anode (Positive): The non-metal anion is oxidised to the pure non-metal.

Example: Electrolysis of Molten Lead(II) Bromide (PbBr₂) (Core 4.1.3a)

Ions present: \(Pb^{2+}\) (cation) and \(Br^-\) (anion)

At the Cathode (Reduction):

• The \(Pb^{2+}\) ions are attracted.
• Reaction: \(Pb^{2+} (l) + 2e^- \longrightarrow Pb (l)\)
• Observation: Droplets of shiny grey molten lead are formed.

At the Anode (Oxidation):

• The \(Br^-\) ions are attracted.
• Reaction: \(2Br^- (l) \longrightarrow Br_2 (g) + 2e^-\)
• Observation: Brown fumes of bromine gas are evolved.

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Section 3: Electrolysis of Aqueous Solutions (The Competition)

When an ionic compound is dissolved in water, the situation becomes more complex because water slightly ionises:

\(H_2O (l) \rightleftharpoons H^+ (aq) + OH^- (aq)\)

This means the electrolyte contains two possible cations (metal ion and \(H^+\)) and two possible anions (salt anion and \(OH^-\)). A competition occurs to see which ion is discharged (reacted) at the electrode.

Prediction Rule at the Cathode (Reduction) (Core 4.1.4)

The product formed depends on the Reactivity Series (Potassium is most reactive; Gold is least).

• Competition: Metal ion vs. \(H^+\) ion.
• Rule: The ion that is easiest to reduce (i.e., the one that is less reactive) will be discharged.

1. Metals Above Hydrogen (K, Na, Ca, Mg, Al): These ions are very stable (hard to reduce).

Result: \(H^+\) ions are discharged to form Hydrogen gas (\(H_2\)).

2. Metals Below Hydrogen (Zn, Fe, Cu, Ag, Au): These ions are less stable (easy to reduce).

Result: The Metal ion is discharged to form a layer of the pure metal.

Analogy: Imagine a VIP queue (the reactive metals). The metal ions (K+, Na+, etc.) are too important/stable to be bothered, so the humble \(H^+\) gets chosen to react instead.

Prediction Rule at the Anode (Oxidation) (Core 4.1.4, Supplement 4.1.10)

The product formed depends on the type of anion and its concentration.

• Competition: Salt anion (e.g., \(Cl^-\), \(SO_4^{2-}\), \(NO_3^-\)) vs. \(OH^-\) ion (from water).
• General Rule: \(OH^-\) is usually discharged to produce Oxygen gas (\(O_2\)), UNLESS one of the halide ions (\(Cl^-\), \(Br^-\), or \(I^-\)) is present at a high concentration.

1. If the Anion is Sulfate (\(SO_4^{2-}\)), Nitrate (\(NO_3^-\)) or Dilute Halide: These anions are very stable and hard to oxidise.

Result: \(OH^-\) ions are discharged to form Oxygen gas (\(O_2\)).

2. If the Anion is a Halide (\(Cl^-\), \(Br^-\), \(I^-\)) and is Concentrated:

Result: The Halogen gas (e.g., \(Cl_2\), \(Br_2\), \(I_2\)) is discharged instead of oxygen.

Common Mistake Alert! Always check the concentration! Dilute NaCl will give oxygen at the anode, but concentrated NaCl gives chlorine gas.

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Section 4: Applying the Rules (Key Examples)

1. Concentrated Aqueous Sodium Chloride (Brine) (Core 4.1.3b)

Ions present: \(Na^+\), \(Cl^-\), \(H^+\), \(OH^-\)

• Cathode (Cations): \(Na^+\) (highly reactive) vs. \(H^+\).
  • Prediction: Hydrogen gas is formed.
  • Reaction: \(2H^+ (aq) + 2e^- \longrightarrow H_2 (g)\)
• Anode (Anions): \(Cl^-\) (concentrated) vs. \(OH^-\).
  • Prediction: Chlorine gas is formed.
  • Reaction: \(2Cl^- (aq) \longrightarrow Cl_2 (g) + 2e^-\)

The overall solution becomes increasingly alkaline because \(Na^+\) and \(OH^-\) ions remain, forming aqueous sodium hydroxide.

2. Dilute Sulfuric Acid (\(H_2SO_4\)) (Core 4.1.3c)

Ions present: \(H^+\), \(SO_4^{2-}\), \(OH^-\) (Since the acid provides plenty of \(H^+\), we only have two ions to compete on the anion side.)

• Cathode (Cations): Only \(H^+\) present.
  • Prediction: Hydrogen gas is formed.
• Anode (Anions): \(SO_4^{2-}\) (stable anion) vs. \(OH^-\).
  • Prediction: Oxygen gas is formed.
  • Reaction: \(4OH^- (aq) \longrightarrow O_2 (g) + 2H_2O (l) + 4e^-\)

3. Aqueous Copper(II) Sulfate (\(CuSO_4\)) Using INERT Electrodes (Carbon/Graphite) (Supplement 4.1.9)

Ions present: \(Cu^{2+}\), \(SO_4^{2-}\), \(H^+\), \(OH^-\)

• Cathode: \(Cu^{2+}\) (unreactive) vs. \(H^+\).
  • Prediction: Copper metal is formed.
  • Observation: A pink/brown solid copper layer forms on the cathode.
• Anode: \(SO_4^{2-}\) (stable anion) vs. \(OH^-\).
  • Prediction: Oxygen gas is formed.
  • Observation: Bubbles of colourless gas are seen.

Did you know? As the \(Cu^{2+}\) ions are removed and replaced by \(H^+\) ions (from the breakdown of water), the solution eventually turns colourless and becomes more acidic (\(H^+\) and \(SO_4^{2-}\) remain, forming \(H_2SO_4\)).

Section 5: Using Reactive Electrodes (The Copper Example)

Up until now, we have assumed the electrodes were inert (unreactive, usually Carbon or Platinum). However, sometimes the electrode itself takes part in the reaction.

Aqueous Copper(II) Sulfate (\(CuSO_4\)) Using COPPER Electrodes (Supplement 4.1.9)

Ions present: \(Cu^{2+}\), \(SO_4^{2-}\), \(H^+\), \(OH^-\)

• Cathode: \(Cu^{2+}\) (unreactive) vs. \(H^+\).
  • Prediction: Copper metal is formed. (The cathode gains mass.)
• Anode (Reactive): Instead of the anions (\(SO_4^{2-}\) or \(OH^-\)) reacting, the copper anode itself is oxidised.
  • Prediction: The copper anode dissolves.
  • Reaction: \(Cu (s) \longrightarrow Cu^{2+} (aq) + 2e^-\) (Oxidation of the electrode)
  • Observation: The copper anode loses mass and gets smaller.

The result: Copper is simply moved from the anode to the cathode. The concentration of the copper sulfate electrolyte remains constant (since copper ions are removed at the cathode and immediately replaced at the anode). This process is used for purifying copper.

Section 6: Writing Ionic Half-Equations (Supplement 4.1.11)

Half-equations show exactly what happens at each electrode in terms of electron loss (oxidation) or gain (reduction).

At the Cathode (Reduction - Gain of Electrons)

1. Hydrogen gas production:

\(2H^+ (aq) + 2e^- \longrightarrow H_2 (g)\)

2. Metal production (e.g., Copper):

\(Cu^{2+} (aq) + 2e^- \longrightarrow Cu (s)\)

At the Anode (Oxidation - Loss of Electrons)

3. Halogen gas production (e.g., Chlorine):

\(2Cl^- (aq) \longrightarrow Cl_2 (g) + 2e^-\)

4. Oxygen gas production (from \(OH^-\) in water/solution):

\(4OH^- (aq) \longrightarrow O_2 (g) + 2H_2O (l) + 4e^-\)

5. Reactive electrode dissolving (e.g., Copper):

\(Cu (s) \longrightarrow Cu^{2+} (aq) + 2e^-\)

Don't worry about memorising the number of electrons for the oxygen equation (4e-), just remember that the reduction reaction at the cathode must be balanced with the oxidation reaction at the anode if you write the overall equation!

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Section 7: Applications of Electrolysis: Electroplating

Electroplating is the process of coating one metal object with a thin layer of another metal using electrolysis.

Reasons for Electroplating (Core 4.1.6)

• To improve appearance (e.g., plating silver onto cheap jewellery).
• To improve resistance to corrosion (e.g., chrome plating steel car parts).

How Electroplating Works (Core 4.1.7)

To successfully electroplate an object (say, a steel spoon with silver), you must set up the cell components correctly:

1. The Object to be Plated: Must be the Cathode (the negative electrode). This is where the metal coating is deposited through reduction.
2. The Plating Metal: Must be the Anode (the positive electrode). This keeps the electrolyte concentration constant (as seen in Section 5).
3. The Electrolyte: Must be an aqueous solution containing ions of the plating metal (e.g., silver nitrate solution if plating with silver).

Example: Silver Plating a Spoon

• Cathode (Spoon): \(Ag^+ (aq) + e^- \longrightarrow Ag (s)\) (The spoon gains silver mass)
• Anode (Silver block): \(Ag (s) \longrightarrow Ag^+ (aq) + e^-\) (The anode replaces the silver ions used)

This setup ensures the silver coating is smooth, continuous, and the process can run for a long time.