Electrochemical Cells SS2 Chemistry Lesson Note

ELECTROLYTIC AND ELECTROCHEMICAL CELL

An electrolytic cell is a device in which electrical energy is converted to chemical energy. An electrochemical cell is a device in which chemical energy is converted to electrical energy.

DIFFERENCES BETWEEN ELECTROLYTIC CELLS AND ELECTROCHEMICAL CELL

Note: An electrolytic cell is also known as a galvanic cell it consists of two half cells:

1. an oxidation half cell
2. a complementary reduction half-cell.

DRAWING AND WRITING OF CELL DIAGRAMS

A typical electrochemical cell consists of two electrodes (cathode and anode) and their electrolytes. The electrodes are connected using a salt bridge. A typical electrode or half-cell is a device in which an element (metal) is in contact with its ions. A typical metallic electrode is symbolically represented as:

M(s)/Mn+ (aq) or Mn+ (aq)/M(s) for metals dipped into solutions of their ions.

Conventional notation is used to describe an electrochemical cell without drawing. Example:

X/X+//+/Y

The single vertical lines (/) represent metal and metal ion interphase while the double lines (//) represent the salt bridge or a porous partition.

For a simple electrochemical cell, there is a flow of electrons from the anode (negative electrode) to the cathode (positive electrode). Also, for an electrochemical to produce electricity, the anode element must be more reactive than the cathode element. That is, the anode element must be higher in the electrochemical series than the cathode.

SALT BRIDGE

In an electrochemical cell, instead of using the porous partition to maintain the electrical neutrality of the half-cell electrolytes, the half-cells can be linked by a SALT BRIDGE. A salt bridge is a filter paper soaked in sodium chloride solution or KCI or NH4NO3.

Functions Of A Salt Bridge

1. It helps in the completion of the electric circuit
2. It enables the migration of ions from one-half cell to the other to maintain electrical neutrality in the solution when current flows in the setup

ELECTRODE POTENTIAL

This is the potential difference that is set up between a metallic electrode and its electrolyte. Example:

1. Copper Ion/copper System:

If a copper plate is dipped into a CuSO4 solution, the net reaction is Cu2+(aq) + 2e- That is: Cu2+(aq)/Cu(s) Cu(s)

2. Zinc Ion/zinc System:

If a zinc plate is dipped into a ZnSO4 solution; the net reaction is:

Zn2+(aq) + 2e > Zn(s)

That is: Zn2+(aq)/Zn(s)

STANDARD ELECTRODE POTENTIAL

The electrode potential of a given system depends on:

1. The overall energy change
2. The concentration of ions in the solution
3. The temperature

Standard electrode potential (Eº) is used to compare the electrode potential values of different metal ions/metal systems.

In measuring the standard electrode potential of metal ion/metal system difference (in volts) which exists between the metallic electrode and the standard hydrogen electrode, the standard hydrogen electrode potential is considered to have an arbitrary value of zero.

E.M.F OF A CELL

When two half-cells are joined together through a salt bridge, the e.m.f. of the cell thus formed is the algebraic difference between the two electrode potentials.

The e.m.f. of the cell formed by the system Zn(s)/Zn2+(aq) & Cu²+(aq)/Cu(s) is defined as the standard electrode potential of the right-hand electrode minus the standard electrode potential of the left-hand electrode.

Etotal = EOR-EL = +0.34-(-0.76) =+1.10V

In general, when this convention is used, then a positive e.m.f indicates that the reaction is

thermodynamically feasible as written down from left to right.

A negative e.m.f implies that the reaction is thermodynamically impossible as written down.