Position of Hydrogen in the Reactivity Series of Metals

Position of Hydrogen in the Reactivity Series of Metals
Reactivity Series
K, Na, Ca, Mg, Al, C, Zn, H, Fe, Sn, Pb, Cu, Hg, Ag, Au
<——– increase in reactivity

Metal oxide + Hydrogen –> Metal + Water
Any metal below hydrogen in the reactivity series, hydrogen will reduce the oxide of metal to metal.
Example 1:

• CuO(s) + H₂(g) –> Cu(s) + H₂O(l)
• Observation: Burns quickly with a bright flame. The black solid turns brown solid.
• H₂: Reducing agent
• CuO: Oxidising agent
• Hydrogen is more reactive than copper.

Example 2:

• PbO(s) + H₂(g) –> Pb(s) + H₂O(l)
• Observation: Burns with a bright flame. The yellow solid turns grey solid.
• H₂: Reducing agent
• PbO: Oxidising agent
• Hydrogen is more reactive than lead.

Example 3:

• Fe₂O₃(s) + 3H₂(g) –> 2Fe(s) + 3H₂O(l)
• Observation: Glows dimly. The reddish-brown solid turns grey solid.
• H₂: Reducing agent
• Fe₂O₃: Oxidising agent
• Hydrogen is more reactive than iron.

Example 4:

• ZnO(s) + H₂(g) –> no reaction
• Observation: No glow is observed. It turns yellow when hot and white when cold.
• Hydrogen is unable to reduce zinc oxide. Hydrogen is less reactive than zinc.

 Redox Reactions in Electrolytic Cell and Chemical Cell
Similarities

• redox reaction.
• Anode: oxidation
• Cathode: reduction
• Electrons flow from anode to cathode in the external circuit

Differences

Differences	Electrolytic Cell (Electrolysis)	Chemical Cell / Voltaic Cell
Structure	With electrical supply.	No electrical supply.
Electrodes	Can be the same or difference metal (graphite or platinum).	Must be two different metals.
Flows of electrons	From anode to cathode through external circuit.	From more electropositive metal to less electropositive metal through external circuit.
Transformation of energy	Electrical energy to chemical energy.	Chemical energy to electrical energy.
At positive terminal	Anode.Oxidation occurs. Anions release electrons at the anode.	Cathode.Reduction occurs. Oxidising agent gain electrons.
At negative terminal	Cathode.Reduction occurs. Cations gain electrons from the cathode.	Anode.Oxidation occurs. Reducing agent releases electrons.

1) Redox Reactions in Electrolytic Cell
Example 1: Electrolysis of molten zinc chloride

• Electrodes: Carbon
• Ions present: Cl^- and Zn²⁺
• Anode: Oxidation / 2Cl^-(l) –> Cl₂(g) + 2e / Cl^- ions act as reducing agent.
• Cathode: Reduction / Zn²⁺(l) + 2e –> Zn(s) / Zn²⁺ ions act as oxidising agent.

Example 2: Electrolysis of copper(II) sulphate solution.

• Electrodes: Carbon
• Ions present: Cu²⁺, SO₄^2-, H⁺, OH^-
• OH^- ions are discharged because OH^- ion is below SO₄^2- ion in electrochemistry series.
  Anode: Oxidation / 4OH^-(aq) –> O₂(g) + 2H₂O(l) + 4e / Oxygen gas is liberated.
• Cu²⁺ ions are discharged because Cu²⁺ ion is below H⁺ ion in electrochemistry series.
  Cathode: Reduction / Cu²⁺(aq) + 2e –> Cu(s) / Cu²⁺ ions are reduced to copper metal (brown layer formed).
• –> Overall equation: Cu²⁺(aq) + 4OH^-(aq) –> O₂(g) + 2H₂O(l) + Cu(s)

Example 3: Electrolysis of copper(II) sulphate solution.

• Electrodes: Copper
• Ions present: Cu²⁺, SO₄^2-, H⁺, OH^-
• OH^- ions  and SO₄^2- ion are not discharged.
  Anode: Oxidation / Cu(s) –> Cu²⁺(aq) + 2e / Copper anode (electrode) dissolves.
• Cu²⁺ ions are discharged because Cu²⁺ ion is below H⁺ ion in electrochemistry series.
  Cathode: Reduction / Cu²⁺(aq) + 2e –> Cu(s) / Cu²⁺ ions are reduced to copper metal.

Example 4: Electrolysis of concentrated sodium chloride solution.

• Electrodes: Carbon
• Ions present: Na⁺, Cl^-, H⁺, OH^-
• Cl^- ions are discharged because of the higher concentration. (Concentration of Cl^- ion is high, the ion is selectively discharged rather than the OH^- ion, the one that is placed below the electrochemical series.)
  Anode: Oxidation / 2Cl^-(aq) –> Cl₂(g) + 2e / Chlorine gas (green gas with choking smell) is liberated.
• H⁺ ions are discharged because Na⁺ ion is below H⁺ ion in electrochemistry series. (H⁺ ions and Na⁺ ion are placed very far apart in the electrochemical series, the concentration factor becomes unimportant.)
  Cathode: Reduction / 2H⁺(aq) + 2e –> H₂(g) / Hydrogen gas is liberated.
• –> Overall equation: 2Cl^-(aq) + 2H⁺(aq) –> Cl₂(g) + H₂(g)

2) Redox Reactions in Chemical Cell
Example 1: Daniel cell

• Anode (negative terminal): Oxidation / Zinc strip immerses in zinc sulphate solution.
  Zn(s) –> Zn²⁺(aq) + 2e / Zinc strip becomes thinner.
• Cathode (positive terminal): Reduction / Copper strip immerses in copper(II) sulphate solution.
  Cu²⁺(aq) + 2e –> Cu(s) / A brown layer formed around copper strip. / Concentration Cu²⁺ ions decreases cause the intensity blue colour of solution decreases.
• Zinc is more electropositive than copper. Electrons are flowed from zinc strip to copper strip through the external circuit. (Note: Conventionally, electrons flow in the opposite direction of electrical current).
• –> Overall equation:  Zn(s) + Cu²⁺(aq) –> Zn²⁺(aq) + Cu(s)