Electrochemical cells gives a current when two half cells are connected with each other. Electrochemical cells is a lesson under electrochemistry section in physical chemistry. Electromotive force (emf) is a another important concept in electrochemistry and it is calculated by half cells in electrochemical cell. Several examples and problems are also explained in this tutorial.
Following sections of electrochemistry are covered in this tutorial.
A simple electrode equilibrium system is made by a metal immersed in an aqueous solution of its salt. As an example, you can think a piece of zinc metal is immersed in water. In the aqueous solution, there is Zn2+ ions such as ZnSO4.
Generally, we can consider metal as M for illustrating our lesson. The metal should process its valance n and the electrolyte must be a solution of Mn+ ions. This means, metal M will remove n number of electrons to form it's Mn+ ions
The metal M has the tendency to from its ion through oxidation (removing electron). And this could be identified as the forward reaction of this system.
M(s) → M2+(aq) + ne
The metal ions (Mn+ ions) in the solution can get accumulated at the metal surface through reduction (accepting electrons). This is the backward reaction of the system.
M2+(aq) → M(s) + ne
Initially (at t=0), the rate of forward reaction is maximum while the rate of the backward reaction is minimum.
Forward reaction proceeds continuously with a decreasing reaction rate and the backward reaction takes place with an increasing reaction rate. After several minutes/ hours the above system could reach the following dynamic equilibrium. (when forward reaction rate and backward reaction rates are equal.)
Any metal electrode made by immersing a metal plate inside a solution of its ion of 1 mol dm-3 at 298K is known as a standard metal electrode. The standard electrode potential could be measured using a voltmeter connected to the metal electrode and the standard hydrogen electrode.
Standard electrode potential of any metal is defined as the potential difference between a metal in equilibrium with its metal ions
( 1mol dm-3) at 298K and the standard hydrogen electrode.
Electrode potentials of metals above H2 are negative while the electrode potentials of metals below H2 are positive.
An electrode made by immersing a platinized platinum plate which absorbed H2(g) at 1 atm pressure and 298K in a solution of 1 mol dm-3 H+ ions known as the standard hydrogen electrode.
This electrode also processes its own absolute electrode potential at standard temperature and pressure. However in electrochemistry we assign this potential as zero for our easiness. And we use this as the standard electrode to measure the electrode potentials of other metal or non-metal electrodes.
Any gaseous electrode could be made through this method by immersing a gaseous electrode in a solution of its ions.
Zn2+(aq) → Zn(s) + 2e
Zn2+(aq) / Zn(s), E = -0.76V
You may see, electrode potential of zinc is a negative value. That means, zinc is oxidized readily to Zn2+ ions. Like that, we can say Zn2+ ions are not readily reduced to Zn due to negative value of electrode potential.
Let's write down the equilibrium reaction between Zn and Zn2+ ions.
Zn2+(aq) ⇌ Zn(s) + 2e
Equilibrium point of this reversible reaction is shifted to left side.
Cl2(g) + 2e → 2Cl-(aq)
Cl2(g) / Cl-(aq), E = 1.36V
You may see, electrode potential of chlorine is a positive value. That means, chlorine is very like to reduce to chloride ions (Cl-). Like that, we can say Cl- is not readily oxidized to Cl2 due to high positive value of electrode potential of chlorine.
Let's write down the equilibrium reaction between Cl2 and Cl- ions.
Cl2(g) + 2e ⇌ 2Cl-(aq)
Equilibrium point of this reversible reaction is shifted to right side.