It is common to separate Redox reactions into its two component parts which is a convenient way of indicating which species gains electrons or loses electrons. Thgese component parts are called HALF REACTIONS. The electrons symbolized in these equations are cancelled out when te half reactions are combined.

e.g. Consider the reaction when zinc is immersed in copper(II) sulphate solution. The overall chemical reaction is expressed as

Zn + CuSO₄ ZnSO₄ + Cu

The reaction can be further expressed by the equation

Zn + Cu ²⁺ Zn ²⁺ + Cu

Zn is being oxidised to Zn ²⁺

Cu ²⁺ is being reduced to Cu

This reaction can be resolved into two half reactions. The one depicting oxidation of zinc

Zn Zn ²⁺ + 2 e ^–

And the other indicating the reduction of copper(II) ions

Cu ²⁺ + 2 e ^– Cu

The driving force for the above reaction can be measured by placing a voltage measuring device in the circuit. We may consider this force as being the sum of two potentials called half cell potentials or single electrode potentials, one of these is associated with the half cell reaction occurring at the anode and the other is associated with the half cell reaction talking place at the cathode.

We cannot measure absolute potentials for half cell reactions but the relative half cell potentials that can be measured are quite useful. To obtain consistent relative half cell potential data, it is necessary to compare all electrodes against a common reference. The reference electrode should be easy to construct, exhibit reversible behaviour, and give constant and reproducable potentials for a given set of experimental conditions. The standard hydrogen electrode (S.H.E.) meets these requirements and is universally used as the ultimate reference electrode. The electrode basically consists of a platinum wire immersed in a solution containing hydrogen ions, and hydrogen gas is bubbled across the surface of the platinum. This type of electrode is called a gas electrode because the platinum takes no part in the electrochemical reaction. The half cell reaction of the cell is given as

H₂ _(g) 2H⁺ + 2 e ^–

By definition this reaction is said to have a zero potential or an E° of zero or E° = 0. By convention, when quoting E° values, the half reactions are written as reductions.

Electrochemical series

is a list of half reactions written in descending order according to their E° values

Half Reaction Oxidizer + electrons reducer	Standard Electrode Potential (volts)
F₂ + 2e ^– 2F^–	+2.87
H₂O₂ + 2H⁺ + 2e ^– 2H₂0	+1.77
Mn0₄^– + 8H⁺ + 5e ^– Mn ²⁺ + 4H₂0	+1.52
Cr₂0₇^2– + 14H⁺ + 6e^– 2Cr3+ 3+ + 7H₂0	+1.36
Cl₂ + 2e ^– 2Cl^–	+1.36
Mn0₂ + 4H+ + 2e ^– Mn ²⁺ + 2 H₂0	+1.28
Br₂ + 2e ^– 2Br^–	+1.07
HN0₂ + H⁺ + e ^– NO + H₂0	+0.99
NO₃^–+ 3 H⁺ + 2e ^– HN0₂ + H₂0	+0.94
Hg ²⁺ + 2e ^– Hg	+0.85
Ag+ + e ^– Ag	+0.80
Fe ³⁺ + e ^– Fe ²⁺	+0.77
0₂ + 2 H⁺ + 2e ^– H₂O₂	+0.68
I₂ + 2e ^– 2I^– +	+0.54
Cu²⁺ + 2e ^– Cu	+0.35
S0₄^2– + 4 H⁺ + 2e^– H₂S0₃ + H₂0	+0.20
S + 2H⁺ + 2e ^– H₂S	+0.14
2H⁺ + 2e ^– H₂	0.00 assigned
Pb²⁺ + 2e ^– Pb	–0.12
Sn ²⁺ + 2e ^– Sn	–0.14
Ni ²⁺ + 2e ^– Ni	–0.25
Fe ²⁺+ 2e ^– Fe	–0.44
2C0₂ + 2 H⁺ + 2e ^– (COOH)₂	–0.49
Cr ³⁺ + 3e ^– Cr	–0.71
Zn ²⁺ + 2e ^– Zn	–0.76
Mn ²⁺ + 2e ^– Mn	–1.05
Al ³⁺ + 3e ^– Al	–1.67
Mg ²⁺ + 2e ^– Mg	–2.34
Na⁺ + e ^– Na	–2.71
Ca ^{2 +} + 2e ^– Ca	–2.87
Ba ²⁺ + 2e ^– Ba	–2.90
K+ + e ^– K	–2.92

–

Using some degree of caution, this electrochemical series can be used to predict which chemical reaction will take place when two chemicals are mixed.

So if we consider the reaction of zinc immersed in copper sulphate again

From the table

The Cu²⁺ / Cu electrode has a value of +0.35V

And the Zn²⁺ / Zn electrode has a value of –0.76V

It can be said, using these values, that the Cu²⁺ will reduce to Cu and the Zn will oxidize to Zn²⁺

(14) Given a diagram or description electrolytic or galvanic cell:

A cell consists of a pair of conductors or electrodes, usually metallic, each of which is immersed in an electrolyte. When the electrodes are connected by an external conductor and a dlow of electrons occurs, a chemical oxidation occurs at the surface of one electrode and a reduction occurs at the surface of the other.

Well a cell is operated to produce electric energy, it is called a Galvanic or sometimes Voltaic cell. A cell requiring an external source of electric energy is called an electrolytic cell

Referring to the diagram of The Electrochemical cell above. This cell is a galvanic cell because when the two electrodes are connected by a wire, electric energy is produced, and a flow of electrons from the zinc electrode to the copper occurs.

This cell can operate as an electrolytic cell if a battery is introduced into the external circuit, which would force electrons to flow in the opposite direction through the cell. In this case Zinc would deposit and copper would dissolve, consuming energy from the battery.

By definition

The Anode is the electrode at which oxidation occurs in the both the electrolytic and galvanic cells. It has a positive charge

The Cathode is the electrode at which reduction occurs in the both the electrolytic and galvanic cells. It has a negative charge .

label the anode and cathode;

So in the diagram,

Oxidation is occurring at the Zinc electrode so this is called the Anode

Reduction is occurring at the Copper electrode so this is called the Cathode

show the direction of the electron flow in the external circuit;

The electrons flow from anode to cathode

;

Ion Movement in the Electrolyte

We have seen that the electrode reactions occurring in the galvanic cell studied above are

Zn Zn ²⁺ + 2 e ^–

Cu ²⁺ + 2 e ^– Cu

As a result, the solution in the cell compartment containing the zinc electrode shows an increase in the concentration of zinc ions, while the solution surrounding the copper electrode is depleted of copper (II) ions. If no interna; contact existed between the two parts of the cell, a charge imbalance would develop in the neighbourhood of the electrodes, e.g. an excess of positive ions would be found near the zinc electrode and an excess of negative ions near the copper electrode.

As a result of this charge imbalance, no current would flow. With the arrangement shown in the diagram, direct contact between the metallic zinc and copper (II) ions is prevented by the porous barrier, but it does not permit the passage of ions.

As a result, zinc ions as well as other cations (positive ions), can migrate from the solution surrounding the zinc electrode toward the copper electrode. Anions (negative ions) can also migrate, but in the opposite direction, that is anions can migrate from the solution surrounding the copper electrode toward the zinc electrode.

The passage of current through a cell involves the migration of ions within the solution, and the current may be considered to be carried by these ions. Not only the ions that react at the electrodes but all ions present in the solution, participate in the carrying of the current,

predict what will happen at the electrodes.

The following points should be noted

1) Electrons always flow from the anode to the cathode

2) Oxidation always occurs at the Anode

3) Reduction always occurs at the Cathode

Reactions of Metals with Water

Sodium and Cold Water

Na Na⁺ + e ^– X2 oxidation

2H₂O + 2 e ^– 2OH^– + H₂ reduction

----------------------------------------------

2Na + 2H₂O 2 Na⁺ + 2OH^–+ H₂

reducer oxidiser

Potassium and Cold Water

K Na⁺ + e ^– X2 oxidation

2H₂O + 2 e ^– 2OH^– + H₂ reduction

----------------------------------------------

2K + 2H₂O 2 K⁺ + 2OH^–+ H₂

reducer oxidiser

Calcium and Cold Water

Ca Ca²⁺ + 2e ^– oxidation

2H₂O + 2 e ^– 2OH^– + H₂ reduction

----------------------------------------------

Ca + 2H₂O Ca²⁺ + 2OH^–+ H₂

reducer oxidiser

Magnesium and Water

Similar reaction to calcium but very much slower in cold water. With hot water, the reaction is somewhat faster but is still quite slow.

If reacted with steam, the magnesium burns forming a white powdery residue.

Mg + H₂O MgO + H₂

reducer oxidiser

Aluminium and Water

Aluminium does not react with water. However, if mercury is rubbed over the surface of the aluminium, the metal slowly displaces hydrogen even from cold water. The lack of reactivity of aluminium can be attributed to the thin layer of oxide that forms on the surface of aluminium.

Zinc and Superheated Steam

If stream is passed over zinc in a furnace to 400°C, a reaction will take place slowly

Zn + H₂O ZnO + H₂

Iron and Superheated Steam

Similar to zinc but a temperature of about 700°C is required

3Fe + 4 H₂O Fe₃O₄ + 4H₂

Fe₃O₄ = (Fe²⁺) (Fe³⁺)₂ (O^2–)₄

A reversible sign is used in the reaction equation because the reverse reaction can occur. So hydrogen is if Hydrogen is passed over hot Fe₃O₄, the oxide is reduced to iron and the hydrogen is oxidised to steam.

Tin, Lead, Copper, Mercury and Silver

These metals do not react with water or steam even at fairly high temperatures.

Reactions of Metals with Hydrochloric Acid

Magnesium

If magnesium is placed in dilute hydrochloric acid it dissolves rapidly with vigorous evolution of a colourless gas which is hydrogen

Mg Mg²⁺ + 2e ^– oxidation

2H⁺ + 2 e ^– H₂ reduction

----------------------------------------------

Mg + 2H⁺ Mg²⁺ + H₂

reducer oxidiser

Two chloride ions will remain unused for each magnesium ion formed. These are spectator ions and partner the magnesium ion in solution. So the solution contains magnesium chloride.

Mg + 2H⁺ + 2Cl^2– Mg²⁺ + 2Cl^2–+ H₂

Aluminium

At first there appears to be no reaction but after a while, particularly if the mixture is warmed, the aluminium begins to react. The delay to the start of the reaction can be attributed to the protective layer of oxide on the surface of the aluminium.

Al Al³⁺ + 3e ^– X 2 oxidation

2H⁺ + 2 e ^– H₂ X 3 reduction

----------------------------------------------

2Al + 6H⁺ 2Al³⁺ + 3H₂

reducer oxidiser

Zinc

The reaction is similar to the reaction of magnesium except it is slower. The hydrogen is produced at a slower rate without frothing and as a result is often used to prepare hydrogen

Zn Zn²⁺ + 3e ^– oxidation

2H⁺ + 2 e ^– H₂ reduction

----------------------------------------------

Zn + 2H⁺ Zn²⁺ + H₂

reducer oxidiser

Iron

Iron is similar to zinc but the reaction is slower and often the acid must be heated to produce a reasonably rapid evolution of gas

Fe Fe²⁺ + 3e ^– oxidation

2H⁺ + 2 e ^– H₂ reduction

----------------------------------------------

Fe + 2H⁺ Fe²⁺ + H₂

reducer oxidiser

Tin

Granulated Tin dissolves very slowly in in cold dilute hydrochloric acid. If the mixture is heated., or if hot concentrated hydrochloric acid is used, the tin dissolves more rapidly

Sn Sn²⁺ + 3e ^– oxidation

2H⁺ + 2 e ^– H₂ reduction

----------------------------------------------

Sn + 2H⁺ Sn²⁺ + H₂

reducer oxidiser

Lead

Lead slowly becomes coated with a white layer of insoluble lead (II) chloride when placed in cold hydrochloric acid, which slows down the reaction. However, lead dissolves fairly rapidly in hot concentrated HCl.

Copper, Mercury and Silver

There is no apparent reaction when these metals are p[laced in hydrochloric acid. Thus they do not reduce the hydrogen ions in aqueous solution of HCl.

The Activity Series

The sequence of metals arranged in order of the readiness with which they react with water or steam is the same as their sequence of reactivity with dilute hydrochloric acid or sulphuric acid and salt solutions. The similarity in sequence suggests a similarity in there actions of the metals in each case. This has been illustrated by showing that each reaction can be interpreted as donations of electrons by atoms of the metals during formation of positive ions. The sequence is called the activity series or the displacement series of the metals. It is also the order of ease of formation of positive ions in solution

Metal	Activity with Cold Water	Activity with Steam	Activity with Acids	Activity with Solutions of Metallic Salts
K	Displace
Na	Hydrogen from
Ca	Cold water
Mg	Do not displace	Displace	Displace	Displace
Al	hydrogen	Hydrogen	Hydrogen from	lower
Zn	From cold water	from steam	Hydrochloric and	Metals
Fe			Dilute sulphuric	from
Sn	Do not displace	Hydrogen	acids	solutions
Pb	From cold water	or steam		of their
Cu	Do not displace	Hydrogen from		salts
Hg	cold water, steam	or acids
Ag

Electroplating

Electroplating Metals With Copper

This is widely employed to forma coating of copper over a base metal, prior to plating with other metals, such as nickel or chromium.

The article to be copper plated is first cleaned of rust or grease and is then made the cathode (negative) in an electrolytic cell. The electrolyte is copper (II) sulphate solution containing a little sulphuric acid. The sulphuric acid improves the conductivity of the electrolyte solution and prevents rough coatings. If firm deposits are to be formed, careful attention must be given to the concentration of the electrolyte, the current used, the size of the cathode and the temperature of the bath.

Production and Purification of Substances

Zinc

The main source of zinc, is zinc blende, which is found mixed with large amounts of other minerals, such as lead sulphide (PbS), pyrites (Fe S₂) and silica (SiO₂). The ore is finely ground and the useful components are separated out by selective flotation. The zinc sulfide is then burnt in air to form zinc oxide and sulfur dioxide

2ZnS + 3O₂ 2ZnO + 2SO₂

The zinc oxide can be reduced to zinc in one of two ways

i) reduction by carbon

A mixture of zinc oxide and carbon (e.g. coke) is strongly heated

ZnO + C Zn + CO

reduction by electrolysis

The zinc oxide is dissolved in dilute sulphuric acid

ZnO + 2H⁺ Zn ²⁺ + H₂O

Zinc dust is then added to the solution to displace all metals below zinc in the activity series. These may be present from impurities in the ore.

e.g. Zn + Cu ²⁺ Zn ²⁺ + Cu _(s)

The solution is electrolysed using a lead anode (which is not attacked by oxygen or acids) and an aluminium cathode. The electrode reactions are

Cathode Reaction

Zn ²⁺ + 2e^– Zn

Anode Reaction

2H₂O 4H⁺ + O₂ + 4e^–

So the zinc ions are replaced by hydrogen ions and the solution contains sulphuric acid. This is used to dissolve the next batch of zinc oxide. The zinc is peeled from the aluminium cathodes when the deposit has reached a suitable thickness

Aluminium

The main source of aluminium is bauxite (Al₂O₃ . x H₂O). The first stage of the process is the purification of the aluminium oxide, which is done by dissolving the bauxite in hot concentrated caustic soda solution under pressure

Al₂O₃ + 3H₂O + 2OH^– 2Al(OH)₄^–

The sodium ions partner the aluminate ions formed and the solution formed is a solution of sodium aluminate.

When the solution is cooled, diluted and seeded with freshly precipitated aluminium hydroxide, most of the aluminate changes into aluminium hydroxide, which precipitates in a coarsely crystalline form and is easily filtered.

2Al(OH)₄^– Al(OH)₃+ OH^–

The precipitate is filtered off, and is heated to decompose it to aluminium oxide, a white powder which is also called alumina

Al(OH)₃ Al₂O₃ + 3H₂O

The alumina is dissolved in a molten mineral called cryolite (Na₃AlF₆) producing an electrolytic conductor. Aluminium is formed at the cathode of a complicated cell.

Dry cells

Carbon Zinc Batteries

The most usual shape is a cylinder with a carbon rod down the centre. This is surrounded by a solid mixture of MnO_2(s), NH₄Cl _(s) and C_(s). The mixture is wet with an electrolyte consisting of a solution containing ZnCl₂ and NH₄Cl. Because NH₄⁺ is an acid, the electrolyte solution is acidic. The outer container of the cylinder is made of zinc.

The electricity produced by the cell is generated by chemical reactions

Of the Zn and MnO₂. The Zn dissolves and releases electrons and hence generates a negative charge on the zinc container.

Zn Zn²⁺ + 2e^–

The MnO₂ consumes H⁺ ions from the electrolyte and electrons from the carbon rod. Hence a positive charge is generated on the carbon rod.

MnO₂+ 4 H⁺+ 2e^– Mn²⁺+ 2H₂O

One disadvantage of this arrangement is that the zinc container develops holes as the zinc dissolves and this allows the electrolyte to leak out. An alternative design uses a pressed carbon outer container in a steel can. Several strips of zinc are placed near the centre of the battery. This ensures much more efficient use of the zinc and also makes the battery leak proof.

Alkaline Batteries

These are essentially the same as carbon-zinc batteries but use an alkaline solution of KOH as an electrolyte. Because the electrolyte is not acidic, steel can be used as the positive terminal instead of carbon and the whole steel container can be sealed, making the battery leak proof. The reaction which generate the electricity in an alkaline battery are

Zn + 4OH^– Zn(OH)₄ ^2–+ 2e^–

MnO₂+ 2H₂O + e^– Mn(OH)₃ + 2OH^–

Rechargable Batteries

They are based on compounds of nickel and cadmium. The electrolytes in these batteries are made by heating a layer of powdered nickel with nickel gauze to produce a porous, flexible layer of the metal. The battery is then assembled from three layers

i) a porous nickel sheet soaked in a solution of nickel salt

ii) a sheet of absorbent paper soaked in potassium hydroxide solution

iii) a porous nickel sheet soaked in a solution of cadmium salt

The sheets are rolled into a tight cylinder and sealed in a steel container. This is fully leak proof.

The KOH reacts with the salts and precipitates insoluble hydroxides into the pores of each metal plate; so one plate is filled with Ni(OH)_2(s) and the other with Cd(OH)_2(s) . The cell is then charged by passing an electric current through it. The Ni(OH)_2(s) plate is made the anode and oxidation occurs at this electrode.

Ni(OH)_2(s)+ 2OH^– NiO_2(s)+ 2H₂O + 2e^–

The Cd(OH)_2(s) plate is the cathode and reduction occurs at this electrode

Cd(OH)_2(s)+ 2e^– Cd_(s)+ 2OH^–

The electrolyte is not consumed, no gas is evolved and, because all the nickel and cadmium compounds are solids, they remain trapped in the pores of the electrodes.

When the cell is used to generate electricity, the reverse reactions take place. The NiO₂consumes electrons and so generates a positive charge. The Cd_(s)releases electrons and so generates a negative charge. In soluble hydroxides are reformed in the pores of the electrodes.

The operation of the cell can be represented by the equation

NiO_2(s)+ Cd_(s)+ 2H₂O

DISCHARGING

CHARGING

Ni(OH)_2(s)+ Cd(OH)_2(s)

Several hundreds of cycles of charging and discharging can be achieved giving a long service life to the battery.

Use redox theory to explain

why corrosion of iron takes place

If a clean iron nail is placed in normal tap water it rusts., but if a nail is placed in freshly boiled air-free water in a sealed flask it does not rust. Also if a nail is kept in dry air in a desiccator it too wont rust. Thus both air and water together are necessary for the corrosion of iron.

Experimental observations have shown that if the humidity of the air is less than 50%, no corrosion occurs. If the humidity is above 80% the iron rusts rapidly.

Other experiments have shown that an electrolyte is needed for corrosion to occur. So NaCl solution (i.e. salt water) will accelerate corrosion.

Irregularities in the surface exposed to the by the metal greatly influence the rate of corrosion. Rusting is very much more for steel than for chemically pure iron. It is also more rapid near imperfections in the surface and near areas of strain, e.g. at the point and head of a nail.

Mechanism of Corrosion of Iron

The water film on the surface of the iron is exposed to the atmosphere and will dissolve any soluble substances that may be present. Thus, carbon dioxide in the atmosphere will dissolve forming carbonic acid, and near the sea other electrolytes. Such as carbon dioxide may be dissolved in the water film. These dissolved substances will make the solution an electrolytic conductor.

The presence of electrolytes in the solution could allow a current producing cell to be set up between the iron as one electrode and an area of impurity, such as a small crystal of carbon, as the other electrode.

A cell can be set in the laboratory to examine the behaviour of iron under these conditions

The production of electrons by iron would involve dissolution of the iron

Fe Fe²⁺ + 2e^– oxidation

The electrons are consumed by a reduction process at the carbon electrode and this accounts for the consumption of oxygen gas.

2H₂O + O₂ + 4e^– 4OH^– reduction

So the corrosion of iron is probably due to the transfer of electrons through the metal to areas of impurity

The areas of impurity are areas to which the electrons drift, because electrons are not being produced at these positions. Thus the oxidation of the iron to the iron (II) state occurs with simultaneous formation of hydroxide ions.

The result of obtaining a solution containing these ions can be seen by adding sodium hydroxide to a solution of an iron (II) salt. A green gelatinous precipitated of iron (II) hydroxide forms.

Fe²⁺ + 4OH^– Fe(OH)₂

On standing, the green iron(II) hydroxide is slowly oxidized to brown iron(III) hydroxide where it is in contact with air .

4Fe(OH)₂ + 2H₂O + O₂ 4Fe(OH)₃

iron (III) hydroxide is hydrated iron (III) oxide

b. the reasons why common methods of corrosion prevention are effective.

Prevention of Corrosion

The corrosion of iron can be greatly diminished or prevented in a number of ways.

i) alloying the iron with other elements

ii) using a protective coating

iii) using electrical protection

iv) unreactive metallic coatings

Alloys

If chromium is alloyed with steel the product is called stainless steel because of its resistance to corrosion. The addition of small amounts of molybdenum further improves its resistance. Cast iron alloyed with silicon is very resistant to corrosion but is weak structurally. However it is useful in building chemical reaction vessels and other articles which require resistance to attack by acids.

Protective Coatings

The most effective method of preventing corrosion of underground steel structures is to completely coat them with an impervious substance such as platicized coal / tar enamel.

Plastic coatings on metal structures are widely used to prevent corrosion. Plastics in use include, polythene, polyvinylchloride (PVC), epoxy resins, rubber and synthetic rubber. Steel can also be coated with glass. These materials provide a barrier to corrosion under even the most severe conditions.

Vitreous enamels and vinyl paints are used to protect car bodies, refrigerators and washing machines.

Electrical Protection

Steel can be protected by being connected to a metal higher than it on the activity series (or electrochemical series). Steel coated with zinc is called galvanised iron and this method of protection is called galvanic protection, or because the steel is protected by preferential corrosion of the zinc, sacrificial protection. Even if the layer of zinc has small imperfections the steel is still protected. The zinc dissolves in preference to the iron because it forms positive ions more readily.

Zn Zn²⁺ + 2e^–

The electrons are consumed on the iron, so preventing corrosion.

Instead of using sacrificial protection, electrons can be supplied by a D.C. generator. This method is used to protect steel wharves. The steel wharf is wired to the negative terminal and the positive terminal is wired to an anode. Only 4 to 5 volts is necessary and a current od a few milliamps per square foot of structure is needed. The anode may be made up of large lumps of steel such as engine blocks. In this case the scrap steel is slowly dissolved.

Fe Fe²⁺ + 2e^–

The anode corrodes away where the electrical connection is made. Silicon – iron anodes are better v=because they corrode only very slowly. The latest development is to use anodes made of titanium coated with a very thin layer of platinum, in which case there is no corrosion of the anodes.

Unreactive Metallic Coatings

Steel is often protected from corrosion by using a thin layer of tin. This is widely used in the manufacture of food containers, (the tin is often coated with an impervious layer of lacquer). With tin plate it is essential that the layers of tin has no imperfections, because tin is below iron on the activity series (above iron on the electrochemical series).

If there is a break in the tin plating, however small, corrosion will commence on the exposed iron and will very rapid. The iron dissolves because it forms positive ions more readily than tin.

Fe Fe²⁺ + 2e^–

The electrons are consumed on the tin and so the iron corrodes. For special purposes, steel is sometimes coated with gold or platinum and it is essential that the layer has no imperfections because these metals are very low on the activity series.

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