1. Chemistry

LESSON 303.10 Electrochemistry

Contents

Introduction4Concepts of Oxidation, Reduction & Redox Reactions4Oxidation4Reduction5Redox Reactions5Electrochemical Cells5OVERALL LESSON SUMMARY18

At the end of this Lesson you should be able to:

describe the basic concepts of Oxidation and Reduction

explain Electrochemical cells

describe Electrochemical potentials

describe the application of Electrochemistry

describe corrosion processes, galvanic corrosion and corrosion prevention methods

Lesson Aims & Objectives

Introduction

Electrochemistry is the study of interchange between chemical energy and

electrical energy. Many significant chemical reactions are electrochemical in

nature. Electrochemical reactions are chemical reactions in which electrons are

transferred. To understand electrochemical reactions, it is necessary to understand

the terms and concepts of electricity and extend these to apply to electrochemical

relationships.

Concepts of Oxidation, Reduction & Redox Reactions

In some reactions, electrons can be lost by atoms and gained by other atoms. This involves oxidation and reduction reactions taking place simultaneously in a process or reaction known as a redox reaction.

Oxidation is the loss of electrons, whereas reduction refers tothe acquisition of electrons, as illustrated in the respective reactions below. The species being oxidized is also known as the reducing agent or reductant, and the species being reduced is called the oxidizing agent or oxidant. The following acronym is useful in remembering this concept:

OIL RIG:

OxidationIsLosing electrons;ReductionIsGaining electrons

Oxidation

Atoms of the element(s) involved in the reaction lose electrons. The charge on these atoms must then become more positive.

Eg: Fe2+ (aq) Fe3+(aq) + e-

Reduction

In reduction reactions, atoms of the elements involved gain electrons.

Eg: Zn2+(aq) + 2e- Zn(s)

Redox Reactions

A redox reaction is an electrochemical reaction in which both reduction and oxidation take place together. The electrons lost in an oxidation component are gained in a reduction component, as shown by an overall equation which is derived from combining the two half reaction equations and cancelling the electrons from both sides.

Eg: Fe3++ Cu+ Fe2+ + Cu 2+ all ions are (aq)

which comes from combining these two half equations:

Fe3+(aq) + e-1 Fe2+(aq) (Reduction)

and

Cu+ (aq) Cu2+ (aq) + e-1 (Oxidation)

Electrochemical Cells

Electrochemical cells contain different components, including;

Ion is an atom or molecule that has an electrical charge.

Cation: An ion that carries a positive charge is called a cation.

Anion: An ion that carries a negative charge is called an anion.

Electrolyte: A solution that contains ions is called an electrolyte solution, or an electrolyte. They conduct electricity as charged ions can move through them.

Astrong electrolyteis a solute that completely, or almost completely, ionizes or dissociates (breaks up into ions) in a solution. These ions are good conductors of electric current in the solution.Strongacids,strongbases, and soluble ionic salts that are not weak acids or weak bases arestrong electrolytes.

Electrodes: A solid electric conductor through which an electric current enters or leaves an electrolytic cell or other medium. In electrochemical reactions oxidation or reduction reaction takes place at the electrodes. These are called electrode reactions, or half-reactions.

A half-reaction can be either a reaction in which electrons appear as products (oxidation) or a reaction in which electrons appear as reactants (reduction). A combination of two half reaction forms the complete reaction.

Cathode: An electrode in which reduction takes place.

Anode: An electrode where oxidation takes place.

Salt Bridge: Adevice used in electrochemistry, to connect theoxidation and reductionhalf-cellsof agalvanic cell(voltaic cell), a type ofelectrochemical cell. It maintains electrical neutrality within the internal circuit, preventing the cell from rapidly running its reaction to equilibrium. If no salt bridge were present, the solution in one half cell would accumulate negative charge and the solution in the other half cell would accumulate positive charge as the reaction proceeded, quickly preventing further reaction, and hence production of electricity.

There are two main types of electrochemical cell, namely the electrolytic cell or the Galvanic cell. Electrons flow around both cells, but the electrolytic cell relies on some form of power supply, as shown in figure 1.

Figure 1 Schematic of Galvanic and Electrolytic Cells

Electrolysis

Ionic substancescontaincharged particlescalledions. For example, lead bromide contains positively charged lead ions and negatively charged bromide ions.

Electrolysisis the process by which ionic substances are decomposed (broken down) into simpler substances when an electric current is passed through them.

For electrolysis to work, the ions must be free to move. Ions are free to move when an ionic substance is dissolved in water or when melted. For example, if electricity is passed throughmoltenlead bromide, the lead bromide is broken down to form lead and bromine.

During electrolysis, positively charged ions move to the negative

electrodeduring electrolysis. They receiveelectronsand arereduced.

Negatively charged ions move to the positive electrode during electrolysis. They lose electrons and areoxidised.

The substance that is broken down is called theelectrolyte. A schematic diagram is shown in figure 2.

Figure 2 Schematic diagram of an Electrolytic Cell

Electrolysis of Simple Electrolytes

Electrolysis of Fused (Melted) Sodium Chloride

An idealized cell for the electrolysis of sodium chloride is shown in the figure below. A source of direct current is connected to a pair of inert electrodes immersed in molten sodium chloride. Because the salt has been heated until it melts, the Na+ions flow toward the negative electrode and the Cl-ions flow toward the positive electrode.

Figure 3 Electrolysis of Fused Sodium Chloride

When Na+ions collide with the negative electrode, the battery carries a large enough potential to force these ions to pick up electrons to form sodium metal.

Cathode:2Na+ (aq) + 2e- 2Na (s)

Anode:2Cl- (aq) Cl2 (g) + 2e-

Overall:2Na+ (aq) + 2Cl- (aq) 2Na(s) + Cl2 (g)

Cl-ions that collide with the positive electrode are oxidized to Cl2gas, which bubbles off at this electrode.

The net effect of passing an electric current through the molten salt in this cell is to decompose sodium chloride into its elements, sodium metal and chlorine gas.

Electrolysis of Aqueous Sodium Chloride

The figure below shows an idealized drawing of a cell in which an aqueous solution of sodium chloride is electrolysed.

Figure 4 Electrolysis of Aqueous Sodium Chloride

Once again, the Na+ions migrate toward the negative electrode and the Cl-ions migrate toward the positive electrode. But, now there are two substances that can be reduced at the cathode: Na+ions and water molecules.

Because it is much easier to reduce water than Na+ions, the only product formed at the cathode is hydrogen gas. Theoretically O2 gas could be produced at the anode, but in practice only Chlorine gas is given off.

Cathode:2H2O (l) + 2e- H2 (g) + 2OH-

Anode:2Cl- (aq) Cl2 (g) + 2e-

Overall:2H2O (l) + 2Cl- (aq) H2 (g) + Cl2 (g) + 2OH- (aq)

Electrolysis of Acidified Water

Electrolysis of water containing sulphuric acid produces oxygen gas at the anode and hydrogen gas at the cathode. The negative sulfate ions (SO42-) or the traces of hydroxide ions (OH-) are attracted to the positive electrode. But the sulfate ion is too stable and nothing happens (it is not discharged). Instead either hydroxide ions or water molecules are discharged and oxidised to form oxygen.

Figure 5 Electrolysis of Acidified Water

The half reactions and overall reaction are as follows:

Cathode:2H2O (l) 4H+ (aq) + O2 (g) + 4e-

Anode:4H+ (aq) + 4e- 2H2 (g)

Overall:2H2O (l) 2H2 (g) + O2 (g)

Electrolysis of Copper Sulphate Solution

An overview and summary of the electrolysis of copper sulphate is shown in figure 6.

Figure 6 Overview and Summary showing Electrolysis of Copper Sulphate Solution

Industrial Applications of ElectrolysisChlor-Alkali Process

The Chlor-alkali process is used to produce several products, namely chlorine gas, hydrogen gas and an alkali solution of sodium hydroxide. These products are all made from the electrolysis of brine (sodium chloride solution or salt water).

Figure 7 shows a summary of the process.

Figure 7 Schematic of the Chlor-Alkali Process

The threeproductsof the electrolysis of concentrated sodium chloride solution have important uses in the chemical industry:

Hydrogen is used as a fuel and for making ammonia

Chlorine is used to kill bacteria in water, and to make bleach and plastics

Sodium hydroxide is used to make soap and bleach

Extraction/Refining of Metals

A block of impure metal is made the anode of an electrolytic cell containing an aqueous solution of the metal salt

A thin sheet of pure metal is made the cathode of the electrolytic cell

When electric current of a suitable voltage is passed, metal ions from the electrolyte get deposited on the cathode as pure metal Mn++ ne- M

Metal ions from the anode enter the electrolyte M Mn++ ne-

Impurities present in the anode settle down as anode mud under the anode

Anode finally disintegrates while the cathode gains in weight due to the collection of pure metal. Figure 8 shows a simple schematic of how pure copper is deposited at the anode.

Figure 8 Diagram of Copper Purification via Electrolysis

Anodising

Anodising is a process that can be used to improve corrosion resistance of metals, a good example being Aluminium and the formation of a protective oxide layer on its surface. This process is illustrated in figure 9.

Figure 9 Process of Anodising Aluminium.

Electroplating

Electrolysis is used to electroplate objects. This is useful for coating a cheaper metal with a more expensive one, such as copper or silver or tin.

Thenegative electrodeshould be the object that is to be electroplated

Thepositive electrodeshould be the metal that you want to coat the object with

Theelectrolyteshould be a solution of the coating metal, such as its metal nitrate or sulphate

Figure 10 Electroplating of Tin

Electropolishing

Electropolishing(and electrochemical polishing) is a production process which removes material. Metal is removed anodically using electrolytes that are specially adapted to suit the material concerned. The aims of electropolishing include a reduction to the roughness of the surface, in other wordsdeburring, together withsmoothness and shine.

Figure 11 summarises a typical process and the key principles of electropolishing.

Figure 11 Summary of Electropolishing Principles

The electrolytes used (chemicals) vary depending upon the metals to be processed. Electropolishing can be applied to workpieces made of aluminium, stainless steel, cobalt alloys, carbon steels, copper and copper alloys, magnesium, nickel and nickel alloys, titanium, zinc, zircon and special metals.

Faradays First Law

Faradays empirical laws of electrolysis relate the current of an electrochemical reaction to the number of moles of the element being reacted and the number of moles of electrons involved.

Faradays Law: the amount of a substance produced or consumed in an electrolysis reaction is directly proportional to the quantity of electricity that flows through the circuit.

The quantity of electricity is related to the current and the time for which it

flows. This is represented by the equation:

quantity of electricity (Q) = current (I) xtime (t)

where, quantity of electricity(Q) is in coulombs

current (I) is in amperes and time (t) is in seconds

In electrolysis calculations, quantity of electricity is measured using the

faraday. The faraday unit has a value of 96 500 coulombs and the number of

electrons in 96 500 coulombs is 6.02 x 1023.

So, 1 faraday = 96 500 coulombs 6.02 1023 electrons = 1 mole of electrons

Faraday's Second Law

The number of faradays of charge which discharges one mole of ions of an

element at an electrode equals the number of charges on the ion.

For example,

1 Faraday will discharge one mole of single charged ions, M+ + 1e- M (s)

2 Faraday will discharge one mole of double charged ions, M2+ + 2e- M (s)

3 Faraday will discharge one mole of triple charged ions, M3+ + 3e- M (s)

(all ion states are aq)

Supposing that the charge required for such reaction was one electron per ion, as is the case for Silver, in the reaction:

Ag+ (aq) + 1e- Ag (s)

According to Faradays law, the reaction with 1 mol of silver would require 1 mol of electrons, or 1 Avogadros number of electrons (6.022 1023). The charge carried by 1 mol of electrons is known as 1 faraday (F). The faraday is related to other electrical units through the electronic charge; the electronic charge is 1.6 1019 coulomb (C). Multiplying the electronic charge by the Avogadro number means that 1 F equals 96,485 C/(mol of electrons).

The amount of electricity that flows, Q, in depositing n moles of Silver, would be calculated using:

Q = nF

Application of Faradays Laws to calculate mass of copper deposited when an electric current is passed through a solution of copper sulphate

Calculate the number of moles of copper and the mass of copper

deposited on the cathode when an electric current of 0.7 A is passed for

45 minutes through a solution of copper sulphate using copper electrodes.

Relative atomic mass Cu = 63.5

1 faraday = 96 500 C

Galvanic Cells

The Daniell cell was the first truly practical and reliable electric battery that supported many nineteenth-century electrical innovations such as the telegraph. In the process of the reaction, electrons can be transferred from the corroding zinc to the copper through an electrically conducting path as a useful electric current. Zinc more readily loses electrons than copper, so placing zinc and copper metal in solutions of their salts can cause electrons to flow through an external wire which leads from the zinc to the copper.

A typical galvanic cell, alongside its typical components, is shown in figure 2.

Figure 2 A typical Galvanic Cell and its Components

Reactivity Series of Metals

The reactivity series of metals can be used to predict which metals will be oxidised and reduced in an electrochemical cell. Metals higher up in the series undergo oxidation more easily and are classed as more reactive than the metals below.

This is useful in predicting which metals are most reactive and preferentially oxidised in a galvanic cell and deposited at the cathode.

Self-assessment Questions

Now complete the following self-assessment questions at the end of this lesson to test your knowledge and application skills.

OVERALL LESSON SUMMARY

Electrolysis

decomposition of compound using electricity

Electrolyte

an ionic compound which conducts electric current in molten or aqueous solution, being decomposed in the process.

Electrode

a rod or plate where electricity enters or leaves electrolyte during electrolysis.Reactions occur at electrodes.

Discharge

the removal of electrons from negative ions to form atoms or the gain of electrons of positive ions to become atoms.

Anode

positive electrode connected to positive terminal of d.c. source.

Oxidation occurs here.

Anode loses negative charge as electrons flow towards the battery, leaving anode positively charged.

This causes anion to discharge its electrons here to replace lost electrons and also, negative charge are attracted to positive charge.

Cathode

negative electrode connected to negative terminal of d.c. source.

Reduction occurs here.

Cathode gains negative charge as electrons flow from the battery towards the cathode, making cathode negatively charged.

This causes cation to be attracted and gains electrons to be an atom.

Anion

negative ion

attracted to anode.

Cation

positive ion

attracted to cathode.

Non-electrolytes

Weak electrolytes

Strong electrolytes

Organic liquids or solutions

Weak acids and alkalis

Strong acids, alkalis and salt solutions

ethanol C2H5OHtetrachloromethane CCl4trichloromethane CHCl3pure water H2Osugar solution C12H22O11molten sulphur S

limewater Ca(OH)2ammonia solution NH3aqueous ethanonoic acid CH3COOHaqueous sulphurous acid H2SO3aqueous carbonic acid H2CO3

aqueous sulphuric acid H2SO4aquous nitric acid HNO3aquous hydrochloric acid HClaqueous potassium hydroxide KOHaqueous sodium hydroxide NAOHcopper(II) sulphate solution CuSO4

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