SSS 2: ELECTRODE POTENTIAL
Electrode Potential is the potential difference set up between a metal and the solution of its salt.
When a metal is dipped into a solution containing its ions, there is a tendency for ions to move between the metal and the solution. This creates an electric potential difference between the metal and its ionic solution called the electrode potential.
The redox potential of a metal is the e.m.f which is measured in volt
The following equilibrium is established when a metal is dipped into a solution containing its ions.
Metal(s) ⇌ Metal ion + electron
M(s) ⇌ M⁺ + e⁻
(plate) (solution) (plate)
Metal ions/ metal system
M⁺(aq) / M(s)
The Copper Electrode System
The copper electrode is represented as:
Cu²⁺(aq) + 2e⁻ ⇌ Cu(s)
(solution) (plate) (plate)
Note:
1. The forward reaction (reduction) represents copper ions gaining electrons to form metallic copper.
2. The backward reaction (oxidation) represents metallic copper losing electrons to form Cu²⁺ ions
3. The positive copper (II) ions from the solution acquire two (2) electrons each from the copper plate to become deposited as a neutral metallic copper atoms on the plate.
4. Copper plate becomes positively charged on the surface due to a deficit of electrons
5. The solution becomes negatively charged due to excess of the negatively charged tetraoxosulphate (VI) ions, SO₄²⁻.
6. A potential difference, known as the electrode potential of copper is set up between the copper metal and the solution containing its ions.
The Zinc Electrode System
(plate) (solution) (plate)
Note:
1. The forward reaction (oxidation) represents metallic zinc losing electrons to form Zn²⁺ ions
2. The backward reaction (reduction) represents Zn²⁺ ions losing electrons to form metallic zinc
3. The neutral zinc atoms on the plate give up electrons to the plate and go into solution as positively charged zinc ions
4. Zinc plate becomes negatively charged on the surface due to excess of electrons
5. The solution becomes positively charged due to excess zinc ions
6. A potential difference known as the electrode potential of zinc is set up between the zinc metal and the solutions containing its ions
Factors Affecting Electrode Potential
The main factors that influence electrode potential include:
1. Concentration of ions:
An increase in the concentration of ions in solution generally increases the electrode potential.
2. Temperature:
An increase in temperature generally decreases the electrode potential. A rise in temperature affects ion mobility and reaction rates, thereby altering the electrode potential.
3. Nature of the metal:
Each metal has a characteristic tendency to lose or gain electrons, which influences its electrode potential.
4. Presence of complexing agents:
Complex-forming substances can change the activity of ions in solution, thereby affecting the electrode potential.
Applications or Uses of Electrode
Potential
1. Electrode potential is used to determine the cell potential and electromotive force (E.M.F.) of an electrochemical cell.
2. It aids in the construction and arrangement of elements in the electrochemical series.
3. It helps to predict the feasibility and direction of redox reactions.
4. It is applied in electroplating processes
5. It is used in the design and development of batteries.
Standard Electrode Potential (E°)
Standard Electrode Potential (E°) of a metal ion/metal system is the potential difference set up between the metal electrode and its ions in solution when the system is in equilibrium under standard conditions.
The standard conditions are:
1. Ion concentration = 1 moldm⁻³ (1 M)
2. Temperature = 25°C (298 K)
3. Pressure = 1 atmosphere (for gases)
4. The system is compared with the Standard Hydrogen Electrode (SHE), which has an assigned potential of 0.00 V.
Factors on which electrode potential
depends
Electrode Potential vary from one metal ions/metal system to another. The electrode Potential of a given system depends on:
1. The overall energy change
2. The concentration of the ions in the solution
3. The temperature
4. The atmospheric pressure
Standard Electrode Potential Value
The absolute value of an electrode potential cannot be measured directly; it must be determined relative to a reference electrode.
The hydrogen electrode has been internationally selected as the standard reference, with its electrode potential arbitrarily assigned a value of zero at all temperatures.
Standard Hydrogen Electrode
2H⁺(aq)/H₂(g)
Note:
The electrode potential value is:
1. Positive: if electrons flow from the hydrogen electrode to the metal electrode
2. Negative: if electron flow from the metal electrode to the hydrogen electrode
Measurement of Standard Electrode
Potentials
1. Measurement of Standard Electrode Potential of Copper
At the Copper electrode
Cu²⁺(aq) + 2e⁻ → Cu(s) [Reduction]
At the platinum electrode
H₂ (g) → 2H⁺(aq) + 2e⁻ [Oxidation]
Overall reaction:
Cu²⁺(aq) + H₂ (g) → Cu(s) + 2H⁺(aq)
Salt Bridge
e.g. a filter paper soaked in sodium chloride acts as a salt bridge.
Salt bridge is represented by a vertical double strokes (||)
Salt bridge contains an inert electrolyte such as:
I. Potassium chloride (KCl)
II.Sodium chloride (NaCl)
III.Sodium sulphate (Na₂SO₄)
IV.Potassium nitrate (KNO₃)
V. Ammonium nitrate (NH₄NO₃)
Functions of salt bridge
Its main functions are:
I. It maintains electrical neutrality by allowing the movement of ions between the half-cells
II. It prevents building up of charges that would stop the cell reaction.
III. It prevents mixing of the different electrolytes in the two half-cells.
IV. It completes the electrical circuit between two half-cells.
Calculations on Standard Electrode
Potential
The e.m.f of the Daniel cell represented by the cell notation:
Zn(s)/Zn²⁺(aq) || Cu²⁺(aq)/Cu(s)
Zn²⁺(aq) + 2e⁻ → Zn(s) : E°= -0.76v
Cu²⁺(aq) + 2e⁻ → Cu(s) : E°= + 0.34v
I. State what the double strokes represent
II. Calculate the standard e.m.f of the set up
Solution:
I. The double strokes represent Salt bridge
II.
= +0.34v - (-0.76v)
= 0.34v + 0.76v
= 1.10v
Example 2:
(a) Calculate the e.m.f of the following reaction:
2Ag(s) + Fe²⁺(aq) → Fe(s) + Ag⁺(aq)
Given that:
Ag⁺(aq)/Ag(s) : E°= + 0.80v
Fe²⁺(aq)/Fe(s) : E°= - 0.44v
(b) Comment on the value of the e.m.f obtained in (a) above
Solution:
2Ag(s) - 2e⁻→ 2Ag⁺(aq).......[Oxidation]
Fe²⁺(aq) +2e⁻→ Fe(s)...........[Reduction]
= -0.44v - 0.80v
= - 1.24v
(b) Since the E° cell is negative, the reaction will not occur spontaneously in the forward reaction
Note:
I. A positive e.m.f means that the system is spontaneous
II. A negative means the reaction is not spontaneous and that the anode cannot reduce the cathode, therefore cell representation should be interchanged
Example 3:
The set up below was used to measure the standard electrode potential of zinc
Zn²⁺(aq) + 2e⁻→ Zn(s)
(b) Show the direction of the electron(s) flow
(c) The e.m.f of the cell above was measured as 1.1v. Taking the standard electrode potential of Cu(s)/Cu²⁺(aq) as +0.34v. Calculate the standard electrode potential of Zn(s)/Zn²⁺(aq) electrode.
Solution:
(a) Zn(s)/Zn²⁺(aq) || Cu²⁺(aq)/Cu(s)
(b) Electrons flow from Zn to Cu
(c)
x = +0.34 - (1.10v)
x = - 0.76v
Note:
(i) The more positive E° value indicates a greater tendency to gain electrons i.e. Reduction occurs (Cathode)
(ii) The more negative E° value indicates a greater tendency to lose electrons i.e. Oxidation occurs (Anode).
(iii) Electrons always flow from anode to cathode through the external circuit
(iv) Conventional current (the direction of positive charge flow) is opposite to electron flow i.e. Current flows from Cu to Zn in the external circuit.
Electrochemical cell
An electrochemical cell (Voltaic or Galvanic cell) is a device that converts chemical energy into electrical energy through a redox reaction.
They are classified into two main types based on whether they can be recharged or not:
1. Primary cells
2. Secondary cells
Primary Cells
Primary cells are electrochemical cells that cannot be recharged once they are discharged.
The chemical reactions in them are irreversible. When the active materials are used up, the cell becomes useless.
Examples include Daniel cell, Dry Leclanché cell (Dry cell) etc
Secondary Cells
Secondary cells are electrochemical cells that can be recharged and reused.
The chemical reactions in them are reversible. The active materials can be regenerated by passing electric current in the opposite direction.
Examples:
I. Lead–acid accumulator (car battery)
II. Lithium–ion (Li–ion) battery etc
Assignment
1. Draw and label the diagram of:
(a) Daniel cell
(b) A Leclanché (dry) cell
(c) A lead–acid accumulator
2. Discuss the charging and recharging of lead–acid accumulator.
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