Electrolysis

Electrolysis is derived from the Greek words meaning "splitting by electricity". An example of electrolysis is the recharging of a car battery. Recharging does not occur spontaneously and must be driven by electrical energy. This is a very important difference between galvanic and electrolytic cells. Reactions taking part during the discharge of a galvanic cell are spontaneous, where as those taking part in an electrolytic cell must be driven by the supply of energy. The difference can be summarised in the diagram on the right.

 

Electrochemical vs electrolytic. Electrical energy is used  in an electrolytic cell to derive chemical energy. This is not a spontaneous reaction. Chemical reactions take part spontaneously to release electrons into an external circuit. Electrical energy can be transformed into chemical energy via an electrolytic cell. Chemical energy can transformed into electrical energy by discharging an electrochemical cell.
Electrolysis of sodium chloride to produce sodium metal and chlorine gas.

Similarities and differences between electrochemical and electrolytic cells are listed in the table below.

Electrochemical (galvanic)
Electrolytic
Oxidation
Occurs at the anode
Occurs at the anode
Reduction
Occurs at the cathode
Occurs at the cathode
Polarity of the anode
Negative
Positive
Polarity of the cathode
Positive
Negative
Energy
Produced
Supplied
Electron flow
From -ve to +ve
From +ve to -ve
Salt bridge Required Not required
Direction of electron flow Anode=>Cathode Anode=>Cathode

 

Lets take the decomposition of molten NaCl(pictured above) as our first example, as this is a relatively simple electrolytic process since only one reaction is possible at each electrode. The power source simply acts as an electron pump pulling electrons out of the anode and pumping them into the cathode.

The chlorine ions are now attracted to the positive electrode (anode) and undergo oxidation.

2Cl-(l) => Cl2(g) + 2e
The sodium ions are attracted to the negative electrode (cathode) and undergo reduction
Na+(l) + e => Na(l)
This process is used in industry to form sodium metal with chlorine as a by-product. The production of sodium takes place in industry in a special electrolytic cell called the Downs cell.

In an electrolytic cell the strongest oxidant and strongest reductant will react.

The strongest reductant will be oxidised at the (+) anode and the strongest oxidant will be reduced at the (-) cathode. Only when these have expired will the next strongest species, according to the electrochemical series react.

Why is NaCl electrolysed in the molten form to produce sodium metal as opposed to the aqueous form? The video on the right shows the electrolysis of a NaCl solution.

What reaction is occurring on the anode?
What reaction is occurring at the cathode?

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The decomposition of molten lead bromide takes place using the apparatus shown on the right, where a current is passed through the electrodes. The reaction proceeds when the switch is closed. Using a table of the electrochemical series answer the following questions.

a) Which electrode is the anode?
b) Which electrode is the cathode?
c) What type of reaction takes place at the anode and the cathode?
d) Write the equation of the reaction taking place at the cathode and the anode. Solution

e) Which electrode would increase in mass? Explain

f) Would a solution of PbBr2 produce the same products? Explain using the table of half cell potential.

Potassium iodide was placed in a u-tube and electrodes, attached to a power source, were inserted at each end. A cotton wool plug was inserted at the bottom of the u-tube, as shown on the left.

The power supply was then turned on and charge was allowed to flow for several seconds. View the video on the left and answer the questions below.

1) What are the possible reactants present?
2) A brown substance is seen to form on the electrode on the right
a) Identify the substance formed
b) Write an equation for the reaction that is occurring at this electrode
c) This electrode is the

3) A gas is seen forming on the electrode on the right. A red litmus is inserted into the solution surrounding the electrode and it quickly turns blue.
a) Identify the gas formed
b) Write the equation for the reaction occurring at this electrode.
c) This electrode is the

4) Name the oxidant and reductant taking part in the half reactions?

5) Write the equation for the overall reaction.

Click for the audio of the solution

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Balance for charge by adding electrons tot he most positive side. Balance for hydrogen by adding Hydrogen ions to the hydrogen deficient side. Balance for oxygen by adding water to the oxygen deficient side. Balance for Ag atoms Balance for oxygen by adding water to the oxygen deficient side. Balance for hydrogen by adding Hydrogen ions to the hydrogen deficient side. Balance for charge by adding electrons tot he most positive side. Home
Each oxygen atom has an oxidation number of -2 Each oxygen atom has an oxidation number of -2 Each potassium atom has an oxidation number of =1