Introduction

  • Electrochemical cells facilitate electron transfer in spontaneous redox reactions between oxidisers and reducers in order to convert energy into a usable form.
  • Two kinds of electrochemical cells: galvanic and electrolytic.
  • Galvanic cells:: Electrochemical cells which convert chemical energy into electrical energy
  • Electrolytic cells:: Electrochemical cells which convert electrical energy into chemical energy.
    • This requires an external energy supply (electricity) to 'force' chemical reactions to occur.
  • Galvanic cells produce electrical energy from redox reactions, and thus, are commonly used in batteries.

Galvanic Cells

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  • Three types of galvanic cells:
    • Metal half cells [voltaic cells with metal (aka 'active') electrodes]
    • Solution half cells [cells with inert (aka 'inactive') electrodes]
    • Fuel cells
  • Consist of two 'half' cells, with each containing an electrode and electrolyte.
  • Electrode:: 'Conductor' with delocalised electrons
  • Electrolyte:: The solution that the electrode is 'dipped' within (the metal ions of the electrode as a soluble salt). This solution conducts electricity.
  • The electrodes are separate, but are connected by a 'salt bridge' (a piece of filter paper dipped in a solution of a salt) - this salt bridge is usually KCl or NaCl.
    • The ions travel through this salt bridge (ions, not electrons)
      • As ions carry charge, having the salt bridge with ions increases the charge carried through.
    • ANOILRIGCAT: Oxidation happens at the anode, reduction happens as the cathode.
  • The galvanometer can also be an ammeter or voltmeter.
  • Electron flow is determined by referring to the metal reactivity series - the more reactive metal in this series.
  • Oxidation only occurs at the anode (-), and the electrons which are lost at this site transfer to the electrode at the cathode across the external (upper) circuit.
  • Reduction only occurs at the cathode (+), and the electrode accepts the transferred electrons (therefore being reduced).
  • Electrons always move from the anode to the cathode.
  • The ions from the two solutions shift to their respective sides; cations to the cathode, and anions to the anode in the internal circuit/salt bridge.

Galvanic Cell Reactions

At the anode in the above image, the following oxidation reaction is occurring:

At the cathode in the above image, the following reduction reaction is occurring:

Over time, there is an eventual formation of pure copper substance at the cathode, and the mass of the zinc at the anode will decrease.

Galvanic Cell Summary Table

Property Anode Cathode
Electron flow Out In
Ion flow Negative (anion) Positive (cation)
Agent Reducing agent Oxidising agent
Charge Negative Positive
Redox Oxidation Reduction

Electrolytic Cells

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  • Electrolysis is the inverse of the redox process, in which electrical energy is converted to chemical energy to power chemical reactions that may not occur otherwise.
  • This process obviously requires electricity to occur (hence why electrolytic cells are so expensive)
  • The anode and cathode are connected via a power supply, with the positive end of the power supply (the lead) connected to the anode.
  • Both the anode and cathode are located in the same container - they share an electrolyte.
  • There is another primary difference between electrolytic cells - the anode is positively charged, and the cathode is negatively charged.
    • Ion flow (as in the electrolyte) does not change.
  • Electrolytic cells are often used to decoratively electroplate cutlery, such as spoons, with a metal.

Electrolytic Cell Reactions

Water electrolysis

  • Electrolysis of water will break down water into its corresponding parts - oxygen gas at the anode, and hydrogen gas at the cathode.
  • Anode:
  • Cathode: [1]

Molten electrolysis

  • Electrolysis of molten will break down salt into its corresponding ions - the chloride ion will be attracted to the anode, and the sodium ion will be attracted to the cathode.
  • This is in the presence of zero water, otherwise the reaction may be impacted.
  • Anode:
  • Cathode:
  1. This is a special case in which hydroxide ions are used to balance the equation, not hydrogen. ↩︎
Interactive Graph
Table Of Contents
Introduction
Galvanic Cells
Galvanic Cell Reactions
Galvanic Cell Summary Table
Electrolytic Cells
Electrolytic Cell Reactions