• Electrochemical cells?generate electricity from spontaneous redox reactions
• For example:

Zn (s)? + CuSO4?(aq)→ Cu (s)? + ZnSO4?(aq)

• • Instead of electrons being transferred directly from the zinc to the copper ions, a cell is built which separates the two redox processes
  • For example:

   

Zn (s)?????Zn2+?(aq)?+ 2e–

• If a rod of metal is dipped into a solution of its own ions, an equilibrium is set up
• Each part of the cell is called a?half?cell

When a metal is dipped into a solution containing its ions an equilibrium is established between the metal and its ions. This is the basis of a half cell in an electrochemical cell.

• This is a?half cell?and the strip of metal is an?electrode
• The position of the equilibrium determines the?potential difference?between the metal strip and the solution of metal
• The Zn atoms on the rod can deposit two electrons on the rod and move into solution as Zn2+?ions:

Zn(s) ? Zn2+(aq) + 2e–

• This process would result in an accumulation of negative charge on the zinc rod
• Alternatively, the Zn2+?ions in solution could accept two electrons from the rod and move onto the rod to become Zn atoms:

Zn2+(aq) + 2e–??? Zn(s)

• This process would result in an accumulation of positive charge on the zinc rod
• In both cases, a?potential difference?is set up between the rod and the solution
• This is known as an?electrode potential
• A similar electrode potential is set up if a copper rod is immersed in a solution containing copper ions (e.g. CuSO4), due to the following processes:

Cu2+(aq) + 2e–??? Cu(s)???– reduction (rod becomes positive)

Cu(s) ? Cu2+(aq) + 2e–??? – oxidation (rod becomes negative)

Oxidation of copper atoms

Reduction of copper(II) ions

• Note that a chemical reaction is not taking place – there is simply a?potential difference?between the rod and the solution
• The potential difference will depend on
  • the?nature of the ions?in solution
  • the?concentration of the ions?in solution
  • the?type of electrode?used
  • the?temperature

   

Electrode potential

• The?electrode (reduction) potential?(E) is a value which shows how easily a substance is?reduced
• These are demonstrated using reversible half equations
  • This is because there is a redox equilibrium between two related species that are in different oxidation states

   

• When writing half equations for this topic, the electrons will always be written on the left-hand side (demonstrating reduction)
• The position of equilibrium is different for different species, which is why different species will have different electrode (reduction) potentials
• The more?positive?(or less negative) an electrode potential, the?more likely?it is for that species to undergo?reduction
  • The equilibrium position lies more to the?right

   

• For example, the positive electrode potential of bromine below, suggests that it is likely to get reduced and form bromide (Br-) ions

Br2(l) + 2e-?? 2Br-(aq)? ? ? ??E?= +1.09 V

• The more?negative?(or less positive) the electrode potential, the less likely it is that reduction of that species will occur
  • The equilibrium position lies more to the?left

   

• For example, the negative electrode potential of sodium suggests that it is unlikely that the sodium (Na+) ions will be reduced to sodium (Na) atoms

Na+(aq) + e-?? Na(s)? ? ? ??E?= -2.71 V

Conventional Representation of Cells

• Chemists use a type of shorthand convention to represent electrochemical cells
• In this convention:
  • A solid vertical (or slanted) line shows a?phase boundary, that is an interface between a solid and a solution
  • A double vertical line (sometimes shown as dashed vertical lines) represents?a salt bridge
    • A?salt bridge has mobile ions that complete the circuit
    • Potassium chloride and potassium nitrate are commonly used to make the salt bridge as chlorides and nitrates are usually soluble
    • This should ensure that no precipitates form which can affect the equilibrium position of the half cells

     

  • The substance with the?highest oxidation state?in each half cell is drawn next to the salt bridge
  • The cell potential difference is shown with the?polarity of the right hand electrode

   

• The cell convention for the zinc and copper cell would be

Zn (s)∣Zn2+?(aq) ∥Cu2+?(aq)∣Cu (s)??????????????????E?cell?= +1.10 V

• This tells us the copper half cell is more positive than the zinc half cell, so that electrons would flow from the zinc to the copper
• The same cell can be written as:

Cu (s)∣Cu2+?(aq) ∥Zn2+?(aq)∣Zn (s)??????????????????E?cell?= -1.10 V

• The polarity of the right hand half cell is negative, so we can still tell that electrons flow from the zinc to the copper half cell

Answer

Al (s)∣Al3+?(aq) ∥ Zn2+?(aq)∣Zn (s)??????????????????E?cell?= +0.94 V

It is also acceptable to include phase boundaries on the outside of cells as well:

∣ Al (s)∣Al3+?(aq) ∥ Zn2+?(aq)∣Zn (s) ∣ ? ? ? ? ? ? ? ???E?cell?= +0.94 V