Electrode Potentials (OCR A Level Chemistry)

Revision Note

Test yourself
Richard

Author

Richard

Last updated

Standard Electrode Potential

Standard electrode potential

  • The position of equilibrium and therefore the electrode potential depends on factors such as:
    • Temperature
    • Pressure of gases
    • Concentration of reagents

  • So, to be able to compare the electrode potentials of different species, they all have to be measured against a common reference or standard
  • Standard conditions also have to be used when comparing electrode potentials
  • These standard conditions are:
    • Ion concentration of 1.00 mol dm-3
    • A temperature of 298 K
    • A pressure of 100 kPa

  • Standard measurements are made using a high resistance voltmeter so that no current flows and the maximum potential difference is achieved

  • The electrode potentials are measured relative to a standard hydrogen electrode
  • The standard hydrogen electrode is given a value of 0.00 V, and all other electrode potentials are compared to this standard
  • This means that the electrode potentials are always referred to as a standard electrode potential (Eθ)
  • The standard electrode potential (Eθis the potential difference ( sometimes called voltage) produced when a standard half-cell is connected to a standard hydrogen cell under standard conditions
  • For example, the standard electrode potential of bromine suggests that relative to the hydrogen half-cell it is more likely to get reduced, as it has a more positive Eθ value

Br2(l) + 2e– ⇌ 2Br(aq)        Eθ = +1.09 V          

2H+(aq) + 2e– ⇌ H2(g)        Eθ = 0.00 V

  • The standard electrode potential of sodium, on the other hand, suggests that relative to the hydrogen half-cell it is less likely to get reduced as it has a more negative Eθ value

Na+ (aq) + e– ⇌ Na(s)        Eθ = -2.71 V

2H(aq) + 2e– ⇌ H2(g)        Eθ = 0.00 V

 

Electrochemical Cells

  • The standard hydrogen electrode is a half-cell used as a reference electrode and consists of:
    • Hydrogen gas in equilibrium with H+ ions of concentration 1.00 mol dm-3 (at 100 kPa)

2H+ (aq) + 2e- ⇌ H2 (g)

    • An inert platinum electrode that is in contact with the hydrogen gas and H+ ions

  • When the standard hydrogen electrode is connected to another half-cell, the standard electrode potential of that half-cell can be read off a high resistance voltmeter

Standard Hydrogen Electrode, downloadable AS & A Level Chemistry revision notes

The standard electrode potential of a half-cell can be determined by connecting it to a standard hydrogen electrode

  • There are three different types of half-cells that can be connected to a standard hydrogen electrode
    • A metal / metal ion half-cell
    • A non-metal / non-metal ion half-cell
    • An ion / ion half-cell (the ions are in different oxidation states)

Metal / metal-ion half-cell

Metal_Metal Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Example of a metal / metal ion half-cell connected to a standard hydrogen electrode

  • An example of a metal/metal ion half-cell is the Ag+/ Ag half-cell
    • Ag is the metal
    • Ag+ is the metal ion

  • This half-cell is connected to a standard hydrogen electrode and the two half-equations are:

Ag+ (aq) + e- ⇌ Ag (s)        E= + 0.80 V

2H+ (aq) + 2e- ⇌ H2 (g)        E= 0.00 V 

  • Since the Ag+/ Ag half-cell has a more positive Evalue, this is the positive pole and the H+/H2 half-cell is the negative pole
  • The standard cell potential (Ecell) is Ecell = (+ 0.80) - (0.00) = + 0.80 V
  • The Ag+ ions are more likely to get reduced than the H+ ions as it has a greater Evalue
    • Reduction occurs at the positive electrode
    • Oxidation occurs at the negative electrode

Non-metal / non-metal ion half-cell

  • In a non-metal / non-metal ion half-cell, platinum wire or foil is used as an electrode to make electrical contact with the solution
    • Like graphite, platinum is inert and does not take part in the reaction
    • The redox equilibrium is established on the platinum surface

  • An example of a non-metal / non-metal ion is the Br/ Br- half-cell
    • Br2 is the non-metal
    • Br- is the non-metal ion

  • The half-cell is connected to a standard hydrogen electrode and the two half-equations are:

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

2H+ (aq) + 2e- ⇌ H2 (g)        E = 0.00 V   

  • The Br/ Br- half-cell is the positive pole and the H/ H2 is the negative pole
  • The Ecellis: Ecell = (+ 1.09) - (0.00) = + 1.09 V
  • The Br2 molecules are more likely to get reduced than H+ as they have a greater Evalue

Non-Metal_Non-Metal Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Example of a non-metal / non-metal ion half-cell connected to a standard hydrogen electrode

Ion / Ion half-cell

  • A platinum electrode is again used to form a half-cell of ions that are in different oxidation states
  • An example of such a half-cell is the MnO4- / Mn2+ half-cell
    • MnO4- is an ion containing Mn with oxidation state +7
    • The Mn2+ ion contains Mn with oxidation state +2

  • This half-cell is connected to a standard hydrogen electrode and the two half-equations are:

MnO4- (aq) + 8H+ (aq) + 5e- ⇌ Mn2+ (aq) + 4H2O (l)       E = +1.52 V

2H+ (aq) + 2e- ⇌ H2 (g)       E= 0.00 V   

  • The H+ ions are also present in the half-cell as they are required to convert MnO4- into Mn2+ ions
  • The MnO4- / Mn2+ half-cell is the positive pole and the H+ / H2 is the negative pole
  • The Ecell is Ecell = (+ 1.52) - (0.00) = + 1.52 V

Ion_ Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Ions in solution half cell

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

Worked example

Writing a cell diagram

If you connect an aluminium electrode to a zinc electrode, the voltmeter reads 0.94V and the aluminium is the negative. Write the conventional cell diagram to the reaction.

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

Examiner Tip

Writing the cell representation is not a specific requirement of the syllabus, however questions will sometimes use cell representations to present information so it is useful to know what a cell representation is.

Students often confuse the redox processes that take place in electrochemical cells.

  • Oxidation takes place at the negative electrode.
  • Reduction takes place at the positive electrode.

Remember, oxidation is the loss of electrons, so you are losing electrons at the negative.

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

You've read 0 of your 10 free revision notes

Unlock more, it's free!

Join the 100,000+ Students that ❤️ Save My Exams

the (exam) results speak for themselves:

Did this page help you?

Richard

Author: Richard

Expertise: Chemistry

Richard has taught Chemistry for over 15 years as well as working as a science tutor, examiner, content creator and author. He wasn’t the greatest at exams and only discovered how to revise in his final year at university. That knowledge made him want to help students learn how to revise, challenge them to think about what they actually know and hopefully succeed; so here he is, happily, at SME.