Sciences - Co-ordinated (Double) (0654) | Electrochemistry

Welcome to Electrochemistry!

Hello future chemist! This chapter, Electrochemistry, connects two huge areas of science: electricity and chemical reactions. Don't worry if it sounds complicated; it’s all about using electrical energy to force non-spontaneous chemical reactions to happen, or using chemical reactions to produce electrical energy. It’s the science behind charging your phone and extracting pure metals!

Let's break down how electricity can drive chemical change—a process called electrolysis.

C4.1 Electrolysis: The Basics (Core & Supplement)

What is Electrolysis? (Core 1)

Electrolysis is the decomposition (breaking down) of an ionic compound, when it is molten (liquid) or in an aqueous solution (dissolved in water), by passing an electric current through it.

For electrolysis to work, we need three key components:

1. Electrolyte: The molten ionic compound or aqueous solution that undergoes decomposition. It must contain free-moving ions.
2. Electrodes: Two rods (usually made of metal or carbon/graphite) that conduct electricity into and out of the electrolyte.
3. Power Source: A battery or power pack providing the direct current (DC) needed to drive the reaction.

Analogy: Think of electrolysis as a tug-of-war. The electricity is the force strong enough to pull apart the ions that are normally stuck together.

Identifying the Electrodes (Core 2)

The electrodes are crucial, and we need to know their names and charges:

• The Anode: This is the positive (+) electrode.
• The Cathode: This is the negative (-) electrode.

Memory Aid: PANC (Positive Anode, Negative Cathode).

Ions are attracted to the electrode with the opposite charge:

• Cations (positive ions) move toward the Cathode (-).
• Anions (negative ions) move toward the Anode (+).

How Charge is Transferred (Supplement 4)

Charge moves in two different ways during electrolysis:

1. In the External Circuit (Wires): Charge is carried by the movement of electrons.
2. In the Electrolyte (Liquid): Charge is carried by the movement of ions (both cations and anions moving in opposite directions).

The actual chemical reactions happen when ions reach the electrodes and either gain or lose electrons:

• At the Anode (+): Anions lose electrons (Oxidation occurs).
• At the Cathode (-): Cations gain electrons (Reduction occurs).

Memory Aid: AN OX, RED CAT (Anode Oxidation, Reduction Cathode).

C4.1 (i) Electrolysis of Molten Compounds

This is the simplest type of electrolysis because only two types of ions are present (the positive metal ion and the negative non-metal ion).

Example: Molten Lead(II) Bromide, \(PbBr_2\) (Core 3a & Supplement 7)

When \(PbBr_2\) is molten, the ions are \(Pb^{2+}\) (cations) and \(Br^-\) (anions).

Step-by-step observations and products:

1. Cations move to the Cathode (-): \(Pb^{2+}\) ions move to the cathode.
2. Anions move to the Anode (+): \(Br^-\) ions move to the anode.
3. Product at Cathode (Reduction): The lead ions gain electrons to form liquid lead metal.
   Observation: Silvery molten metal forms/collects at the negative electrode.
4. Product at Anode (Oxidation): The bromide ions lose electrons to form brown bromine gas.
   Observation: Red-brown fumes of bromine gas are seen leaving the positive electrode.

Writing Half-Equations (Supplement 8)

Half-equations show the electron transfer at a single electrode. The syllabus requires us to construct ionic half-equations for reactions at the cathode (showing gain of electrons as a reduction reaction).

At the Cathode (Reduction):

Lead(II) ions gain 2 electrons to become lead metal:
\[ Pb^{2+}(l) + 2e^- \longrightarrow Pb(l) \]

At the Anode (Oxidation): (For completeness, showing the required oxidation reaction)
Bromide ions lose electrons to form diatomic bromine gas:
\[ 2Br^-(l) \longrightarrow Br_2(g) + 2e^- \]

C4.1 (ii) Electrolysis of Aqueous Solutions

When an ionic compound is dissolved in water, the electrolysis gets trickier because the water itself provides extra ions: \(H^+\) and \(OH^-\). Now there is a competition at each electrode to see which ion reacts!

The General Rules for Aqueous Electrolysis (Supplement 6)

1. At the Cathode (-) [Reduction]

Ions competing: The metal cation (e.g., \(Na^+\)) and the \(H^+\) ion (from water).

Rule: Metals or hydrogen are formed at the cathode.

• If the metal ion is more reactive than hydrogen (e.g., \(Na^+\), \(K^+\)), hydrogen gas is produced.
• If the metal ion is less reactive than hydrogen (e.g., \(Cu^{2+}\), \(Ag^+\)), the metal itself is produced.

2. At the Anode (+) [Oxidation] (Inert Electrodes)

Ions competing: The anion (e.g., \(Cl^-\), \(SO_4^{2-}\)) and the \(OH^-\) ion (from water).

Rule: Non-metals (other than hydrogen) are formed at the anode.

• If a halide ion (\(Cl^-\), \(Br^-\), \(I^-\)) is present, and it is concentrated, the halogen gas is produced (e.g., chlorine gas).
• If the solution is dilute or contains polyatomic ions like sulfate (\(SO_4^{2-}\)) or nitrate (\(NO_3^-\)), then oxygen gas is produced from the \(OH^-\) ions in the water.

Don't worry if this reactivity series concept seems tricky. Just remember that highly reactive metals (like sodium or potassium) are "lazy" and prefer to stay in solution, leaving hydrogen to react instead.

Example A: Concentrated Aqueous Sodium Chloride (\(NaCl\)) (Core 3b)

Electrolyte Ions: \(Na^+\), \(Cl^-\), \(H^+\), \(OH^-\). (Using inert carbon/graphite electrodes).

• At the Cathode (-): \(Na^+\) vs. \(H^+\). Sodium is highly reactive. Hydrogen gas is produced.
  Observation: Bubbles of colourless gas.
  Reduction Half-Equation:
  \[ 2H^+(aq) + 2e^- \longrightarrow H_2(g) \]
• At the Anode (+): \(Cl^-\) vs. \(OH^-\). Since the solution is concentrated, chlorine gas is produced (instead of oxygen).
  Observation: Bubbles of pale green/yellow gas (chlorine).

Example B: Dilute Sulfuric Acid (\(H_2SO_4\)) (Core 3c)

Electrolyte Ions: \(H^+\), \(SO_4^{2-}\), \(H^+\), \(OH^-\). (Using inert platinum/carbon electrodes).

• At the Cathode (-): Only \(H^+\) ions are present (from the acid and water). Hydrogen gas is produced.
  Observation: Bubbles of colourless gas.
• At the Anode (+): \(SO_4^{2-}\) vs. \(OH^-\). Sulfate ions do not react. Oxygen gas is produced from the \(OH^-\) ions (water oxidation).
  Observation: Bubbles of colourless gas.

Example C: Aqueous Copper(II) Sulfate (\(CuSO_4\)) using Inert Electrodes (Supplement 5)

Electrolyte Ions: \(Cu^{2+}\), \(SO_4^{2-}\), \(H^+\), \(OH^-\). (Using inert carbon/graphite electrodes).

• At the Cathode (-): \(Cu^{2+}\) vs. \(H^+\). Copper is less reactive than hydrogen. Copper metal is produced.
  Observation: A pink/brown solid (copper) deposits on the electrode.
  Reduction Half-Equation:
  \[ Cu^{2+}(aq) + 2e^- \longrightarrow Cu(s) \]
• At the Anode (+): \(SO_4^{2-}\) vs. \(OH^-\). Sulfate ions do not react. Oxygen gas is produced.
  Observation: Bubbles of colourless gas.

Did you know? As the \(Cu^{2+}\) is used up and the \(H^+\) (acidic ion) is left behind, the blue colour of the solution gradually fades, and the solution becomes more acidic.

Example D: Aqueous Copper(II) Sulfate (\(CuSO_4\)) using Copper Electrodes (Supplement 5)

When the electrode itself is made of an active metal (like copper), the anode reaction changes.

Electrolyte Ions: \(Cu^{2+}\), \(SO_4^{2-}\), \(H^+\), \(OH^-\). (Using active copper electrodes).

• At the Cathode (-): Same as before: Copper metal deposits. (Purification/Electroplating)
• At the Anode (+): Now, the copper anode itself participates. Instead of \(OH^-\) or \(SO_4^{2-}\) reacting, the copper metal loses electrons (it dissolves).
  Observation: The anode gets smaller and loses mass.
  Oxidation Half-Equation:
  \[ Cu(s) \longrightarrow Cu^{2+}(aq) + 2e^- \]

Key Takeaway: In active electrode electrolysis (like copper refining), the anode dissolves and the cathode gains mass, but the overall concentration of the electrolyte usually stays constant.

C4.2 Hydrogen-Oxygen Fuel Cells (Core & Supplement)

Now we look at the opposite process: using a chemical reaction to generate electricity.

What is a Fuel Cell? (Core 1)

A hydrogen-oxygen fuel cell uses the chemical reaction between hydrogen (\(H_2\)) and oxygen (\(O_2\)) to produce electricity.

The fuel cell is electrochemical cell where the reactants flow in continuously, and the electrical energy is generated continuously.

The overall chemical product of this reaction is just water:
\[ 2H_2(g) + O_2(g) \longrightarrow 2H_2O(l) + \text{Electrical Energy} \]

The major advantage is that it is a clean energy source, producing no greenhouse gases or pollutants, only water!

Comparing Fuel Cells to Traditional Engines (Supplement 2)

Hydrogen-oxygen fuel cells are often proposed as replacements for traditional combustion engines that run on gasoline/petrol. Here is a comparison:

Feature	Hydrogen-Oxygen Fuel Cells	Gasoline/Petrol Engines
Energy Source	Hydrogen (renewable source potential)	Fossil fuels (non-renewable)
Pollution	Only chemical product is water (zero pollution at point of use).	Produces \(\text{CO}_2\), \(\text{NO}_x\), and particulates (causes acid rain, global warming, and respiratory issues).
Efficiency	High energy conversion efficiency.	Lower efficiency (a lot of energy is wasted as heat and noise).
Maintenance	Fewer moving parts, potentially lower maintenance.	Complex, high maintenance required.
Refuelling/Storage	Hydrogen production and safe storage are currently challenging and expensive.	Infrastructure is well established; fuels are easy to store (as liquids).