Galvanic Cell Video

Cell Potential

CELL POTENTIAL (E_cell)

1. Difference in electrical potential of electrodes.
2.  Also called voltage or electromotive force (emf).

• Cell emf depends on the concentration of ions in the cell, the temperature of solutions and pressure of gasses involved.

·         Unit : Volt (V).

Ecell > 0	Ecell = 0	Ecell < 0
Spontaneous cell reaction.	Reaction has reach equilibrium.	Non spontaneous cell reaction.

STANDARD CELL POTENTIAL (E^O_cell)

1. Difference in electrical potential of electrodes measured at specified temperature (usually 298 K) with all components in their standard state which are :

•  1 atm (for gases)
•  1M for solutions
• Pure solid for electrodes 

 2. The E^o_cell can be calculated using the equation:
              E^o_cell = E^o_cathode - E^o_anode

·       STANDARD ELECTRODE (HALF-CELL) POTENTIAL (E^O_half-cell)

1. Potential associated with a given half reaction (electrode compartment) when all components are in their standard states.
2. Also called standard reduction potential.
3. Changing the balancing coefficients of a half-reaction does not change E^o value because electrode potentials are intensive properties.
4. Eg. :

                    Oxidation (anode)     :      Zn(s) → Zn²⁺(aq) + 2e^-          E^o_anode = -0.76 V

                    Reduction (cathode)  :     Cu²⁺(aq) + 2e^- → Cu(s)         E^o_cathode = +0.34 V

                                                                                                                                                                                                                                                                                                                                                                                                                                                                                                                           

                   Overall                      :     Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

E^o_cell = E^o_cathode - E^o_anode

        = E^o_copper - E^o_zinc

        = +0.34 V – (-0.76 V)

        = +1.10 V

STANDARD HYDROGEN ELECTRODE (SHE)

1. SHE is the standard reference electrode to measure the electrode potentials of other half-cell.
2. SHE is assign as zero potential (0 V)SHE consists of H₂ gas at 1 atm,25°C bubbling around a platinum electrode which is immersed in 1 M of solution H⁺.
3.  The electric potential of the half reaction is due to the reduction of H⁺ to H₂ (g).

               2H⁺ (aq) +2e^-→H₂ (g)      E° =0.00V

STANDARD REDUCTION POTENTIAL TABLE

From the table,

1. The element with more +ve E° becomes cathode.
2. The element with more -ve E° becomes anode.
3. The more +ve E^o favours reaction to the right.
4. The more -ve E^o favours reaction to the left.
5. The more +ve E° , the stronger the oxidising agent.
6. The more -ve E° , the stronger the reducing agent.

Cell Notation

•          The Galvanic cell can represented by the cell diagram / cell notation.

Half-Cell Equation

HALF-CELL EQUATION

Oxidation (anode)        :           Zn(s) → Zn²⁺(aq) + 2e^-

Reduction (cathode)    :           Cu²⁺(aq) + 2e^- → Cu(s)

                                                                                                           

Overall                        :          Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Salt Bridge

SALT BRIDGE

• Contains a solution of non reacting ions (KCl, KNO₃, Na₂SO₄).
• It functions is to maintain electrical neutrality while redox reactions proceed at the electrodes.

OXIDATION HALF-CELL (ANODE)	REDUCTION HALF-CELL (CATHODE)
Zn (s) → Zn2+(aq) + 2e-	Cu2+(aq) + 2e- → cu (s)
·         Zn2+ ions enter the solution causing an overall excess of positive charge. ·         Anode loses mass/erode ·         Cl- ions from salt bridge move into anode (Zn) half-cell.	·         Cu2+ ions leave the solution causing an overall excess of negative charge. ·         Cathode gain mass/deposit ·         K+ ions from salt bridge move into cathode (Cu) half-cell.

Galvanic Cell

Comparison: Galvanic Cell and Electrolytic Cell

Electrochemical Cells:

1) Galvanic Cells

2) Electrolytic Cells

Similarities

·        They consist of an electrolyte each. Ø  Electrolyte is mixture of ions (usually in aqueous solution) that are involve in reaction or that carry charge. ·         They consist of two electrodes. Ø  Electrodes are objects that conduct electricity between cell and surroundings. Ø  Electrodes consist of :                                                      I.            Anode – Oxidation half-reaction takes place                                                      II.          Cathode – Reduction half-reaction takes place ·         The process of donation of electrons occurs at the anode while the process of acceptance of electrons occurs at the cathode. ·         Electrons flow from the anode to the cathode in the external circuit.

Galvanic Cells	Differences	Electrolytic Cells
Converts electrical energy to chemical energy	Conversion of energy	Converts chemical energy to electrical energy
a) The anode (positive terminal) is positively charged. b) The cathode (negative terminal) is negatively charged.	Charge of a) Anode b) Cathode	a) The anode (negative terminal) is negatively charged. b) The cathode (positive terminal) is positively charged.
Carbon of two same metals or two different metals.	Type of electrodes	Two different metals.
Spontaneous reaction. -Reaction that has a natural tendency to occur and does not require an energy input for it to occur. -eg : rusting of iron nails	Type of reaction	Non-spontaneous reaction. -Reaction that cannot occur naturally and needs energy input to help it to occur. -eg : photosynthesis

Redox Reaction

REDOX REACTION

• A process in which there is a net movement of electrons from one reactant (reducing agent) to another (oxidizing agent).
• Chemical reaction involving oxidation and reduction occurs simultaneously.
• Also called oxidation-reduction reaction.

COMPARISON BETWEEN OXIDIZING & REDUCING AGENT

OXIDIZING AGENT	REDUCING AGENT
Substance that undergoes reduction	Substance that undergoes oxidation
Substance that accepts electrons in a redox reaction	Substance that loss electrons in a redox reaction
Substance that undergoes a decrease in oxidation number	Substance that undergoes an increase in oxidation number

Electrochemistry

Electrochemistry is a study of relationship between chemical change and electrical work.

OXIDATION	REDUCTION
Loss of electron	Gain of electron
Gain of oxygen	Loss of oxygen
Loss of hydrogen	Gain of hydrogen
Increase in oxidation number	Decrease in oxidation number
Eg.                    : Zn(s) → Zn2+ (aq) + 2e-   Oxidation no.    :   0            +2	Eg.                    : 2H+(aq) + 2e- → H2 (g) Oxidation no.    :   +1                      0