Unlock the Power of Chemical Cells

Dive into the world of chemical cells and discover how they convert energy. Learn about electrodes, electrolytes, and their applications in modern technology. Master the principles of electrochemistry with our engaging lessons.

In this lesson, we will learn:

• To recall the basic function of an electrochemical cell.
• To explain the role of the key features in an electrochemical cell in redox processes.

• We learned in Introduction to electrochemistry that electrochemistry is about chemical potential energy (the energy ‘locked in’ the bonds of all chemical substances) being converted into electrical energy (energy from the movement of electrons).
  
  
• Converting chemical potential to electrical energy requires electrons to move FROM one reactant TO the other reactant – which is where we get the processes of reduction and oxidation from, and the ‘redox’ phrase so common in electrochemistry. Redox processes that produce electricity occur in electrochemical cells, or chemical cells, which work just like electrical circuits (like in physics) but they also use ions in solution to carry charge, not just electrical wiring.
  The ‘battery’ of an electrochemical cell is the energy in the bonds of the reactants that are getting broken up!
  
  
• The two most important features of a chemical cell are the two electrodes where redox happens:
  • The anode is the positively charged electrode where oxidation happens.
    • It attracts the negatively charged ions in the cell.
    • The negatively charged ions lose electrons (they are oxidized) at the anode, and the electrons flow through to the cathode.
  • The cathode is the negatively charged electrode where reduction happens.
    • It attracts the positively charged ions in the cell.
    • The positively charged ions gain electrons (they are reduced) at the anode by the electrons that have been flowing to the anode from the cathode.
  • Chemical cells work by electrons flowing from the anode to the cathode.
  • Electrode is just a general term for either cathode or anode, a conducting material where the redox processes happen. They must conduct electricity to work, so if the reactants are not in the solid state (see electrolysis in a later lesson) they would be made of a metal like platinum, or graphite.
  
  
• Other features in a chemical cell, or redox cell, are:
  • The salt bridge, which is used to connect the two half-cells. It contains a solution of highly soluble spectator ions (e.g. KNO₃ or NaNO₃) to carry current without taking part in the reaction.
  • Wire will connect the two electrodes and allow electrons to flow from the anode to the cathode. This completes the ‘circuit’ (just like a circuit in physics!) for a one-way flow of electrons.
  • A voltmeter can be added to measure a flow of electrons.
  
  
• The real working of a redox cell can be split into the two ‘half-reactions’, the two separate redox processes of reduction and oxidation.
  For example with Ni metal (in NiCl₂ solution) reacting with Pb metal (in Pb(NO₃)₂ solution), the two half-reactions for the metal atoms are:
  • Ni²⁺ + 2e^- $\rightleftharpoons$ Ni
  • Pb²⁺ + 2e^- $\rightleftharpoons$ Pb
  Redox works by one substance being oxidized, losing an electron(s), and one substance being reduced, gaining an electron(s) – the electron(s) that was lost by what got oxidized! So which substance will do which process?
  Together, chemists have studied many redox experiments and with their combined results, a table of ‘standard reduction potentials’ (how easily can this substance be reduced?) was formed – all good chemistry textbooks will have this table in the back. For our example, Pb is more likely to be reduced than Ni,¹ so lead will be reduced and nickel will be oxidized in this cell.
  • Because lead has a greater tendency to be reduced, when the half-cells are connected via wiring and a salt bridge, electrons will flow toward the lead electrode. This upsets the equilibrium between lead metal and Pb²⁺ ions:
    
    Pb²⁺ + 2e^-$\enspace \rightleftharpoons \enspace$Pb
    The increase in electrons available will shift the equilibrium to the right to counter the imbalance, which produces lead metal in a reduction reaction. This makes the lead electrode the cathode of this cell.
  • Nickel has a smaller tendency to reduce than lead does, so it will do the oxidation when these two half-cells are connected up. When connected, electrons will flow away from the nickel electrode and toward the lead electrode through the wire. This disturbs the nickel equilibrium we wrote:
    
    Ni²⁺ + 2e^-$\enspace \rightleftharpoons \enspace$Ni
    The decreasing amount of electrons will shift the equilibrium to the left to counter to produce more electrons. In doing this, it converts nickel metal into Ni²⁺ ions and e^- in an oxidation reaction. This makes the nickel electrode the anode of this cell.
    
    
    ¹ Source for data: ATKINS, P. W., & DE PAULA, J. (2006). Atkins' Physical chemistry. Oxford, Oxford University Press.