Electrochemical Cells

• An electrochemical cell can allow two half-reactions to occur separately, with electrons flowing through an external wire, from one reaction to the other.
• This is the basis in all batteries and “dry” cells.
• In one part of the cell, the oxidation reaction is taking place; in the other part, the reduction reaction is occurring.
• The electrons that have been lost in the oxidation reaction travel through the external circuit to the reduction reaction, where they reduce the oxidising agent.
• The two parts of the electrochemical cell are called half-cells.
• The energy is given out as electrical energy instead of heat.
• Batteries have both a positive and negative terminal; the potential difference between the positive and negative terminals is measured in volts.
• The voltage can also be expressed as the number of joules of energy transferred when one coulomb of charge flows around the circuit:
  1 volt = 1 J C^-1
  
• For example, when 1 coulomb of charge flows through a potential difference of 3 volts, 3J of energy are transferred. After an amount of time, the voltage in a battery can drop; this is due to the current; the higher the current drawn, the lower the voltage the battery may give. Thus, when comparing cells and half-cells, it is important that the potential difference is measured when no current flows. This potential difference is given the symbol E_cell (sometimes referred to as electromotive force, or emf, even though it’s not actually a force).
• A high resistance voltmeter is used to measure the potential difference between the electrodes of the two half-cells. It is high resistance, so that negligible current flows, and therefore the maximum potential difference is being measured.
• From the measured potential differences of half-cells, it is possible to work out a scale of individual electrode potentials.
• This sort of scale can be used to tell us, for any pair of half-cells, which one would be the positive electrode.
• Electrons flow to the half-cell with the more positive potential and cause a reduction reaction; an oxidation reaction must occur in the other half-cell.
• Therefore, using a scale of electrode potentials, it is possible to determine what is oxidised and what is reduced in a Redox reaction.

Metal- metal ion cells

• A simple cell can be established by using strips of metal submersed into solutions of metal ions; for example, the copper and zinc cell consists of two half-cells:
• A connection is needed between the two solutions (the solutions should not mix together).
• A strip of filter paper soaked in saturated potassium nitrate can be used as a junction between the two half-cells (salt bridge). The salt bridge contains both positive and negative ions. Whichever half-cell becomes more negative (i.e. contains the oxidising agent) will receive positive ions from the salt bridge, cancelling the charge. The half-cell that becomes more positive (i.e. the one containing the reducing agent) will receive negative ions from the saltbridge, cancelling the charge.
• In a half-cell, there is an equilibrium reaction occurring at the surface of the electrode; for example in the zinc half-cell:
  Zn_(s) Zn²⁺_(aq) + 2e
  
• When the two half-cells are connected together, the species that is most reactive will have the greatest tendency for it to form ions i.e. become the reducing agent.
• When the zinc half-cell is connected to the copper half-cell; the copper ions will accept electrons, as they are the stronger oxidising agent. This will shift the position of equilibrium in the copper half-reaction to the right, thus resulting in the reduction of copper ions.
  Cu _(s) Cu²⁺_(aq) + 2e
  
• A high resistance voltmeter is connected between the copper and zinc electrodes. The table below shows the data from a number of different cells:
• It is hard to detect any patterns in the voltages; different pairs of half-cells are giving almost the same voltage in some cases.
• One way of sorting out some of the data is show below.
• The potential difference between the copper and iron half-cells is 0.78V. The potential difference between the iron and zinc half-cells is 0.32V (iron the positive half-cell). Therefore the difference between the zinc and copper half-cells must be 1.10V.
• A hydrogen half-cell is used as a standard reference; i.e. the hydrogen half-cell is used to construct a list of electrode potentials.
• Standard conditions are used, so as not to alter the voltage (voltage changes with temperature).
• The potential of the hydrogen half-cell is defined as 0V.
• The standard electrode potential of a half cell, , is defined as the potential difference between it and a standard hydrogen half-cell.
• Some values are shown below:
• The table above is called the electrochemical series.
• The metal with the most negative electrode potential has the highest tendency to act as a reducing agent, supplying electrons when it is oxidised.
• The metal with the least negative electrode potential has a much lower tendency to act as a reducing agent.
• The table can be used to allow us to calculate the maximum voltage obtainable from any standard cell.
• It can also be used to make predictions about any Redox reaction, and whether or not it will occur in a cell.
• The electrochemical series can be found in your data book.

Other half-cell reactions

• In addition to just metal/metal ion reactions, there are also many other kinds of half-cell reactions, for example:

  Between ions:
  

  
  Fe³⁺ (aq) + e^- Fe²⁺ _(aq)
  Cr₂O₇^2- _(aq) + 14H⁺_(aq) + 6e^- 2Cr³⁺ _(aq) + 7H₂O _(l)
  MnO₄^- _(aq) + 8H⁺_(aq) + 5e^- Mn²⁺ _(aq) + 4H₂O _(l)
  Between molecules and ions:
  Cl_2(aq) + 2e^- 2Cl^- _(aq)
  

• With half-cells like these, electrical contact is made with an inert electrode dipping into the solution containing all the ions and molecules involved in the redox half-reaction.

The Electrochemical Series

• The more positive the value of a species in the series, the stronger the oxidising agent.
• From the electrochemical series, you are able to work out in which direction the half-cell reaction will occur.
• The electrode potential is a measure of the tendency of a half-reaction to supply electrons:
  • Zinc Metal in a zinc solution has an electrode potential of -0.76V, and so has a high tendency to supply electrons.
  • Copper metal in a copper solution has an electrode potential of +0.34V, and so has a lower tendency to supply electrons.
  • Silver metal in a silver solution has an electrode potential of +0.80V and so has a very low tendency to supply electrons.
• So, zinc can supply electrons to copper, which can supply electrons to silver, but the reverse cannot occur.
• For zinc to supply electrons, the reaction must be:
  Zn_(s) Zn²⁺_(aq) + 2e^-
  
• For copper to accept electrons, the reaction must be:
  Cu²⁺_(aq) + 2e^- Cu_(s)
  
• The prediction for the overall reaction agrees with the observed changes:
  Zn_(s) + Cu²⁺ Zn²⁺ + Cu_(s)
  
• The highest electrode potential in two connected half-cells is always the positive electrode of a cell. Electrode potential charts provide a useful way of displaying and using the electrochemical series data:

Using Electrode Potential Charts

Example Question:

The table below gives electrode potentials of some half-reactions. What reactions can occur if we connect the MnO₄^-_(aq)/Mn²⁺_(aq) and Fe³⁺_(aq) half-cells, so that electrons can flow? Write an equation for the overall reaction.

Method:

1. Construct an electrode potential chart.
2. Use it to make predictions about half-reactions.
3. Balance the half-equations.

1.  
2. Electrons flow to the positive electrode of a cell; this will be to the MnO₄^- _(aq) /Mn²⁺_(aq) half-cell (this is more positive than the Fe³⁺_(aq) / Fe²⁺_(aq) half-cell.
   
   A reduction reaction will occur at the positive electrode:
   MnO₄^-_(aq) + 8H⁺_(aq) + 5e^- Mn²⁺_(aq) + 4H₂O_(l)
   
   The other half-cell will supply the electrons:
   Fe²⁺_(aq) Fe³⁺_(aq) + e^-
   
3. Now to balance the equations to make sure they match in number of electrons:
   5Fe²⁺_(aq) 5Fe³⁺_(aq) + 5e^-
   The overall reaction will be:
   MnO₄^-_(aq) + 8H⁺_(aq) + 5Fe²⁺_(aq) Mn²⁺_(aq) + 4H₂O_(l) + 5Fe³⁺_(aq)

Using electrode potentials to predict reactions

• Electrode potentials can be used to make predictions about any redox reaction.
• However, these are only predictions, and therefore are not certainty.
• Electrode potentials can only be used to decide whether a reaction is possible; even if it is possible it may be too slow to actually occur.
• If a reaction is slow, it is often because the activation energy is too low.
• So, even if the electrode potentials state that a reaction will occur; in reality the activation energy is too high for it to work.
• We could make the reaction work by adding a catalyst, which lowers the activation energy, or by changing the conditions (i.e. temperature, pressure etc.).
• If the electrode potentials state that a reaction cannot occur, then the reaction will never occur, no matter what is done to try and make it happen.

Useful books for revision