AQA A Level Chemistry复习笔记5.4.2 Standard Electrode Potentials

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 (Eꝋ)
• The standard electrode potential (Eꝋ) is 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 Eꝋ value

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

2H+(aq) + 2e– ⇌ H2(g)        Eꝋ = 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 Eꝋ value

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

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

• 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

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

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)        Eꝋ = + 0.80 V

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

• Since the Ag+/ Ag half-cell has a more positive Eꝋ value, this is the positive pole and the H+/H2 half-cell is the negative pole
• The standard cell potential (Ecellꝋ) is Ecellꝋ = (+ 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 Eꝋ value
  • 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 Br2 / 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)        Eꝋ = +1.09 V

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

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

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)       Eꝋ = +1.52 V

2H+ (aq) + 2e- ⇌ H2 (g)       Eꝋ = 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 Ecellꝋ is Ecellꝋ = (+ 1.09) - (0.00) = + 1.09 V

Ions in solution half cell

Standard cell potential

• Once the Eꝋ of a half-cell is known, the potential difference or voltage or emf of an electrochemical cell made up of any two half-cells can be calculated
  • These could be any half-cells and neither have to be a standard hydrogen electrode

   

• The standard cell potential (Ecellꝋ) can be calculated by subtracting the less positive Eꝋ from the more positive Eꝋ value
  • The half-cell with the more positive Eꝋ value will be the positive pole
    • By convention this is shown on the right hand side in a conventional cell diagram, so is termed Erightꝋ

     

  • The half-cell with the less positive Eꝋ value will be the negative pole
    • By convention this is shown on the left hand side in a conventional cell diagram, so is termed Eleftꝋ

     

   

Ecellꝋ = Erightꝋ - Eleftꝋ

• • Since oxidation is always on the left and reduction on the right, you can also use this version

Ecellꝋ = Ereductionꝋ - Eoxidation

Answer

Step 1: Calculate the standard cell potential. The copper is more positive so must be the right hand side.

Ecellꝋ = Erightꝋ - Eleftꝋ

Ecellꝋ = (+0.34) - (-0.76)

= +1.10 V

The voltmeter will therefore give a value of +1.10 V

Step 2: Determine the positive and negative poles

The Cu2+ / Cu  half-cell is the positive pole as its Eꝋ is more positive than the Eꝋ value of the Zn2+ / Zn half-cell