Electrolysis

Electrolysis A summary 

Simple Electrolytic Cell

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Introduction •

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Electrolysis is the process of decomposing a compound into its constituent elements by electricity. During electrolysis electrolysis::  –

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A direct current is passed into the compounds in its aqueous or molten state Electrical energy is transformed into chemical energy. Ions are usually discharged.

Introduction •

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An electrolyte is a compound which conducts an electric current when molten or dissolved in water. During electrolysis electrolysis,, anions are attracted to the anode and they are oxidised. Cations are attracted to the cathode and they are reduced.

Electrolysis of molten lead bromide

Electrolysis of molten electrolytes

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The metallic ions will be discharged and reduced at the cathode to form a metal. Mn+ + ne-  M The metal will be the cathode product.

Non-metallic ions are oxidised to non-metal at the anode. The non-metal will be the anode product. Example Electrolysis of molten lead (II) bromide Bromide ion is oxidised to bromine gas and the gas is liberated at the anode. 2Br- (aq)  Br2 (g) + 2e Lead (II) ions are reduced to lead metal and the metal is deposited at the cathode. Pb2+ (aq)+ 2e  Pb(s) •

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Electrolysis of solutions •

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At the respective electrodes, there will be a selective discharge of ions. Factors affecting the selective discharge are: position in the Reactivity Series relative concentrations nature of electrodes

Position in the Reactivity Series •

A cation which is lower in the Reactivity series will be preferably discharged than another higher in the series.

Reactivity Series Cations

Anions

Potassium ion, K +

Fluoride ion, F-

Sodium ion, Na+

Sulfate ion. SO42-

Calcium ion, Ca 2+

Nitrate ion, NO3-

Magnesium ion, Mg 2+ Zinc ion, Zn2+

Ease of discharge increases

Chloride ion, Cl Bromide ion, Br-

Iron ion, Fe2+

Iodide ion, I-

Lead ion, Pb2+

Hydroxide ion, OH-

Hydrogen ion, H + Copper(II) ion, Cu2+

Electrolysis of solutions Apparatus :

Electrolysis of aqueous electrolytes Example Electrolysis of dilute sulfuric acid Electrolyte : H2SO4 (aq) Ions in electrolyte : H+, OH- and SO42Electrodes : carbon (graphite) •

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Electrolysis of dilute H2SO4

Another picture of the electrolysis of dilute sulfuric acid

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At the anode (+): OH- and SO42- ions migrate to the anode. Position of OH- is lower than SO42- in the electrochemical series. OH- ion is preferentially discharged. Each OH- ion loses one electron and is oxidised to oxygen and water. 4OH- (aq)  2H2O (g) + O2 (g) + 4e Bubbles of oxygen gas are liberated.

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At the cathode (-): Hydrogen ions migrate to the cathode. H+ ion is discharged. Each H+ ion gains one electron and is reduced to hydrogen gas. 2H+ (aq) + 2e  H2 (g) Bubbles of hydrogen gas are liberated.

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Conclusion: Water is decomposed into hydrogen and oxygen. 2H2O (l)  2H2 (l) + O2 (g) H2SO4 is not decomposed. H2SO4 is added to increase the electrical conductivity of the electrolyte solution. As water is decompos decomposed, ed, the concentration concentratio n of sulfuric acid increases.

Concentration of Ions •

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The ion present in higher concentration tends to crowd out other ions and hence gets discharged at the electrode. The following ions are not  discharged even though they are present at high concentrations: Anions : F-, SO42-. NO3- and CO32Cations : K+. Na+, Ca2+, Mg2+ and Al3+

Electrolysis of concentrated sodium chloride solution •

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Electrolyte : concentrated NaCl (aq) Electrodes : Carbon (graphite) Ions in the electrolyte : H+, Na+, OH-, Cl-

Electrolysis of concentrated NaCl (aq)

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At the anode(+): OH- and Cl- ions migrate to the anode. Concentration of Cl - ions is higher than that of OH-. Cl- ions is preferentially discharged. Each Cl- ions loses one electron and is oxidised to chlorine. 2Cl-  Cl2 + 2e Greenish-yellow chlorine gas is liberated.

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At the cathode (-): H+ and Na+ ions migrate to the cathode. Position of Na+ ion is too high up the electrochemical series to be discharged. Each H+ ion gains one electron and is reduced to hydrogen. 2H+ + 2e  H2 Bubbles of hydrogen gas are liberated.

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Conclusion: Hydrogen and chlorine gases are formed. Sodium hydroxide is left behind in the electrode. 1 volume of hydrogen to 1 volume of chlorine is collected.

Electrolysis in Industry •

Electroplating of metals is carried out to :  –

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Improve the appearance of a metal Prevent corrosion of a metal

The object to be electroplated must be made the cathode cathode.. The anode must be the metal used for plating. The electrolyte solution must contain ions of the metal for plating.

Example •

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To copper plate a table spoon. Electrolyte : CuSO4 Ions : Cu2+, H+, SO42-, OHAnode: pure copper Cathode : table spoon

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At the anode (+) OH- and SO42- ions migrate to the anode The 2 ions are not discharged. Each copper atom loses two electrons and is oxidised to Cu2+. Cu (s)  Cu2+ (aq) + 2e Copper anode dissolves

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At the cathode (-) H+ and Cu2+ ions migrate to the cathode. Cu2+ ion is preferentially discharged. discharged. Each Cu2+ ion gains two electrons and is reduced to copper. Cu2+ (aq) + 2e  Cu(s) A thin layer of copper metal is deposited on the surface of the table spoon.

Another example of electroplating copper •

Copper metal is deposited at the graphite cathode.

Simple Electric Cell •

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A device which converts chemical energy into electrical energy is called a cell, or battery. It consists of a pair of dissimilar metals in an electrolyte.

Simple Electric Cell

A simple cell : using lemon juice as an electrolyte •

Two different metals are used as the electrodes. A voltage is registered in the electronic voltmeter.

Simple Electric Cell •

Of the two metals, the more reactive metal:  –

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Undergoes oxidation Loses electrons more readily Becomes the anode Becomes the negative terminal

The cell voltage depends on the position of the two metals in the reactivity series. The further apart the two metals, the bigger the voltage produced.