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SS2: CHEMISTRY - 1ST TERM

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  1. Periodicity and Periodic Table I | Week 1
    5 Topics
    |
    1 Quiz
  2. Quantum Numbers Orbitals & Electrical Structure | Week 2
    6 Topics
    |
    1 Quiz
  3. Periodicity and Periodic Table II | Week 3
    12 Topics
    |
    1 Quiz
  4. Periodicity and Periodic Properties III | Week 4
    11 Topics
    |
    1 Quiz
  5. Periodicity and Periodic Properties IV | Week 5
    5 Topics
    |
    1 Quiz
  6. Mass-Volume Relationship in Reaction | Week 6
    8 Topics
    |
    1 Quiz
  7. Types of Reactions: Oxidation and Reduction | Week 7 & 8
    7 Topics
    |
    1 Quiz
  8. Oxidation – Reduction Reaction II | Week 9
    3 Topics
    |
    1 Quiz
  9. Electrode Potential and Electrochemical Cells I | Week 10
    6 Topics
    |
    1 Quiz
  10. Electrode Potential and Electrochemical Cells II | Week 11
    5 Topics
    |
    1 Quiz
  11. Electrolysis I | Week 12
    8 Topics
    |
    1 Quiz
  12. Electrolysis II | Week 13
    8 Topics
    |
    1 Quiz
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Topic Content:

  • Measuring Standard Electrode Potential
  • Standard Electrode Potential of Copper Half-cell
  • Standard Electrode Potential of Zinc Half-cell

To measure the standard electrode potential of metal ions, its electrode is connected to a standard hydrogen electrode via a salt bridge and a voltameter to show the readings. A salt bridge is an inverted U-tube containing a saturated solution of KNO3 or KCl, its function is that it allows the migration of ions between two electrodes.

Electrode potentials are measured under the following standard conditions:

(i) concentration of the electrolyte in 1.0 mol dm-3

(ii) gaseous reactant at one atmospheric pressure

(iii) temperature of the cell at 25°C or 298K

Under these conditions, the electrode potential is called standard electrode potential Eo. The superscript (o) stands for the standard state.

Standard Electrode Potential of Copper Half-cell:

When the Cu2(aq)/Cu(s)half-cell is connected to the  2H+(aq)/H2(g) half-cell by a salt bridge the voltameter reads 0.34 Volts. This is the potential difference between the two half cells and it is known as the electromotive force (e.m.f) of the cell.

The half-cell reaction at the copper and hydrogen electrode shows that electrons flow from the hydrogen electrode to the copper electrode.

Reactions at the Electrodes:

At hydrogen electrode:

H2(g) → 2H(aq) + 2e (Oxidation)

At Copper electrode:

Cu2+(aq) + 2e → Cu(s) (Reduction)

Overall reaction:

Cu2+(aq) + H2(g) →  Cu(s) + 2H+(aq)

The positive sign in e.m.f of copper shows that copper ions Cu2+ has a greater attraction for electrons than H

Cu2+(aq) / Cu(s) : Eo = + 0.34 volts

Standard Electrode Potential of Copper Half cell
Cu electrode joined to the H2 electrode

Standard Electrode Potential of Zinc Half-cell:

When Zn(s)/Zn(aq) half-cell is connected to the 2H+(aq) /H2(g) half-cell by a salt bridge, the voltameter reads – 0.76 Volts, but with the zinc electrode forming the negative terminal relative to the standard hydrogen electrode.

The half-cell reaction at the zinc and hydrogen electrode shows that electrons flow from the zinc electrode to the hydrogen electrode. Oxidation takes place at the zinc electrode and reduction at the hydrogen electrode. 

Reactions at the Electrodes:

At the zinc electrode:

Zn(s) →   Zn2+(aq) + 2e (Oxidation)

At the hydrogen electrode: 

2H+(aq) + 2e →  H2(g) (Reduction)

Overall Reaction

Zn(s) +  2H+(aq) →  Zn2+(aq)  + H2(g)

The negative sign in e.m.f of zinc shows zinc has the potential of losing electrons to hydrogen ions or hydrogen ions have a greater attraction for electrons than Zn2+

Zn2+(aq) / Zn(s), Eo = -0.76 Volts

Standard Electrode Potential of Zinc Half cell
Zn electrode joined to the H2 electrode