Oxidation+Reduction.

OXIDATION - REDUCTION
An oxidation-reduction reaction (or redox reaction) is one that involves the transfer of electrons from one species to another.
 * Definition **

Oxidation was originally defined as a reaction with oxygen. Mg reacts with O2 to form magnesium oxide, MgO.  2Mg(s) + O2(g) ®  2MgO(s) Simlarly reduction was the removal of oxygen e.g. CO reduces Fe2O3 and produces Fe and CO2.  Fe2O3(s) + 3CO(g) ®  2Fe(s) + 3CO2(g) There are also reactions that are classed as redox reactions that do not involve oxygen e.g. the displacement by zinc metal of copper metal from a solution of copper sulfate.  Zn(s) + Cu2+(aq) ®  Zn2+ (aq) + Cu(s) In this reaction the Zn metal has been oxidised as it has lost electrons, and the Cu2+ has been reduced as it has gained electrons. Zn ®  Zn2+ + 2e -  and Cu2+ + 2e - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cu

=<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hence the mnemonic - = =<span style="font-family: Arial,sans-serif; font-size: 210%;"> **OIL RIG** = =<span style="font-family: Arial,sans-serif; font-size: 160%;"> **o**xidation **i**s **l**oss **r**eduction **i**s **g**ain =

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The species that is oxidised, in this case Zn, is called the reducing agent or **reductant.** <span style="font-family: Arial,sans-serif; font-size: 11pt;">The species that is reduced, in this case Cu2+, is the oxidising agent or **oxidant.** <span style="font-family: Arial,sans-serif; font-size: 11pt;"> //reduced oxidised// <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3O4 + 4C ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 3Fe + 4CO <span style="font-family: Arial,sans-serif; font-size: 11pt;"> //oxidant reductant// <span style="font-family: Arial,sans-serif; font-size: 11pt;"> e.g. NaCl + AgNO3 ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> NaNO3 + AgCl <span style="font-family: Arial,sans-serif; font-size: 11pt;">Reactions which involve covalent compounds cannot be adequately described in terms of loss and gain of electrons. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> e.g. S + O2 ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> SO2 and PCl3 + Cl2 ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> PCl5 <span style="font-family: Arial,sans-serif; font-size: 11pt;">Rather than consider gain or loss of oxygen or electrons, a better way is to use a system of **oxidation numbers.** In all redox reactions there is a change in oxidation numbers.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Reactions in which there is no electron transfer are NOT redox reactions.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Oxidation numbers **

<span style="display: block; font-family: Arial,sans-serif; font-size: 120%; text-align: center;">An **increase in oxidation number** corresponds to **oxidation** <span style="display: block; font-family: Arial,sans-serif; font-size: 120%; text-align: center;">A **decrease in oxidation** number corresponds to **reduction**. = = =<span style="font-family: Arial,sans-serif; font-size: 11pt;">Oxidation number rules = <span style="font-family: Arial,sans-serif; font-size: 11pt;">The oxidation number (or state) is a number that can be assigned to each atom in an element, <span style="font-family: Arial,sans-serif; font-size: 11pt;">compound or ion, using a set of rules. These rules are as follows: <span style="font-family: Arial,sans-serif; font-size: 11pt;">e.g. in Na2O the oxidation numbers are still +1 for Na+ and -2 for O2 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">The oxidation number of an atom in an element is **zero.** For example in H2 the oxidation number of H is 0.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">The oxidation number of an atom in a monatomic ion is the same as the charge on the ion e.g. in Na+ the oxidation number is +1, in O2 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> the oxidation number is -2. In an ionic compound containing monatomic ions the ions have the same oxidation numbers as they would alone
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">In compounds each hydrogen atom usually has an oxidation number +1 (the exception is in the metal hydrides e.g.NaH where oxidation number of H= -1).
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">In compounds each oxygen atom has oxidation number -2 (except in peroxides e.g. H2O2).
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">In a molecule the sum of the oxidation numbers of all the atoms is zero.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Using rules 3, 4 and 5 it is possible to calculate the oxidation numbers of all atoms. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Find the oxidation number of S in H2SO4. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2 x +1 + 1x ? + 4 x -2 = 0 <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Oxidation number of S = +8 - 2 = **+6** //<span style="font-family: Arial,sans-serif; font-size: 11pt;">Exercise 1 //<span style="font-family: Arial,sans-serif; font-size: 11pt;">: //Find the oxidation numbers of the N atoms in each of the following molecules.// //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> NO2, HNO3, NO, N2, N2O, HNO2, N2O4 //

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> In the ion Cr2O72 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> the oxidation number of Cr is calculated as follows: <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2 x ? + 7 x -2 = -2 <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2 x? = -2 + 14 = +12 <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Oxidation number of Cr = +12 /2 = +6
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">In **polyatomic ions** the sum of the oxidation numbers of all the atoms is equal to the charge on the ion.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Oxidation numbers are always quoted **per** atom.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Balancing Redox Equations. <span style="font-family: Arial,sans-serif; font-size: 11pt;">It is not always possible to balance a reaction equation by trial and error. The following method for balancing redox equations is commonly called the ion-electron half-equation method.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Step 1 **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Identify the oxidation and reduction reactions and the appropriate reactant and product in each case. For the reduction reaction the species reduced will be the reactant with an oxidation number that decreases. The opposite is true for the oxidation step e.g. when a solution of potassium dichromate reacts with iron II nitrate the species oxidised is Fe2+ and the species reduced is Cr2O72 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">.

<span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">and

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Cr2O72 -® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cr3+


 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Step 2 **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Balance all atoms undergoing a change in oxidation number.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">and <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cr2O72 - ® **<span style="font-family: Arial,sans-serif; font-size: 11pt;">2 **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Cr <span style="font-family: Arial,sans-serif;">3+

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">and <span style="font-family: Arial,sans-serif; font-size: 11pt;">Cr2O72 -® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> **2**Cr3+ + **7H2O**
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Step 3 **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Balance the number of O atoms by adding the appropriate number of water molecules.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">and <span style="font-family: Arial,sans-serif; font-size: 11pt;">Cr2O72 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> + **14H** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> **2**Cr3+ + **7H2O** <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Step 4 **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Balance the H atoms by adding H+ ions.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Step 5 **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Balance the charge by adding electrons, e-. This gives 2 balanced half-equations.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ + **e** - <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">and <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cr2O72 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> + **14H+** + **6e** -® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> **2**Cr3+ + **7H2O**
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> In the oxidation half-equation the Fe2+ loses electrons and in the reduction half-equation the Cr2O72 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> gains electrons

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ + **e** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> (x6) => **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> 6Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 6Fe3+ + 6**e** - <span style="font-family: Arial,sans-serif; font-size: 11pt;">Added to Cr2O72 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> + 14H+ + 6e -® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2Cr3+ + 7H2O <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">gives the final equation **<span style="font-family: Arial,sans-serif; font-size: 11pt;">6Fe2+ + Cr2O72 ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + 14H+ ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> **2Cr3+ + 7H2O + 6Fe3+** <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;"> <span style="font-family: Arial,sans-serif; font-size: 11pt;">Finally check that the equation is balanced, particularly for charge!!
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Step 6 **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> To obtain an overall balanced equation the 2 half equations must be added together. Before doing this the equations have to be multiplied so that the number of electrons in each half-equation is the same. In this way, the electrons will be eliminated in the final equation.

=<span style="font-family: Arial,sans-serif; font-size: 11pt;">Common Oxidants = <span style="font-family: Arial,sans-serif; font-size: 11pt;">Both species are colourless so no colour change is observed.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">1. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Oxygen gas, O2 **<span style="font-family: Arial,sans-serif; font-size: 11pt;">, involved in all burning reactions producing the oxide ion, O2 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">2. **Halogens** e.g. chlorine, Cl2 (a yellow-green gas), bromine, Br2 (an orange liquid), and iodine, I2 (a shiny black solid) are all reduced to their respective colourless halide ions, Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;">, Br - <span style="font-family: Arial,sans-serif; font-size: 11pt;">, I - <span style="font-family: Arial,sans-serif; font-size: 11pt;">. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> The order of oxidising strength is Cl2 > Br2 > I2. Any halogen is able to oxidise the halide ion from a weaker halogen e.g. Cl2 can successfully oxidise I - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> to I2. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Cl2 is commonly used as a disinfectant and in swimming pools. It is added as a white solid Ca(OCl)2 rather than as chlorine gas.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">3. **Hydrogen peroxide, H2O2** is a colourless liquid that is reduced to water. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> H2O2 + 2H+ + 2e- ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2H2O

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> The more active the metal the more violent the reaction.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">4. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hydrogen ions, H+ **<span style="font-family: Arial,sans-serif; font-size: 11pt;">present in dilute acids is reduced to hydrogen gas, H2. Only metals above hydrogen in the activity series will react with H+ in dilute acid or water.


 * ~ Activity Series ||~  ||~   ||~   ||
 * = <span style="font-family: Arial,sans-serif; font-size: 160%;">potassium, ||= K ||= Kate's ||  ||
 * = <span style="font-family: Arial,sans-serif; font-size: 160%;">sodium ||= Na ||= Nanny || <span style="font-family: Arial,sans-serif; font-size: 14px;">Most reactive metals react violently with water ||
 * = <span style="font-family: Arial,sans-serif; font-size: 160%;">calcium ||= Ca ||= Can || <span style="font-family: Arial,sans-serif; font-size: 14px;">and are not found naturally in elemental form. ||
 * = <span style="font-family: Arial,sans-serif; font-size: 160%;">magnesium ||= Mg ||= Manage ||  ||
 * = <span style="font-family: Arial,sans-serif; font-size: 160%;">aluminium ||= Al ||= A ||  ||
 * = <span style="font-family: Arial,sans-serif; font-size: 160%;"> zinc ||= Zn ||= Zebra ||  ||
 * = Iron ||= Fe ||= For ||  ||
 * = tin ||= Sn ||= She || <span style="font-family: Arial,sans-serif; font-size: 14px;">React slowly with water and acid ||
 * = lead ||= Pb ||= Punishes ||  ||
 * = //Hydrogen// ||= H ||= Him ||  ||
 * = copper ||= Cu ||= Cruelly ||  ||
 * = mercury ||= Hg ||= How || <span style="font-family: Arial,sans-serif; font-size: 14px;">Metals below H do not react with acid. ||
 * = silver ||= Ag ||= Agonizingly ||  ||
 * = gold ||= Au ||= Awful ||  ||
 * = silver ||= Ag ||= Agonizingly ||  ||
 * = gold ||= Au ||= Awful ||  ||

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Zn(//s//) + Cu2+(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Zn2+(//aq//) + Cu(//s//) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> On the other hand, Ag will not react with copper sulfate solution as Ag is below Cu in the
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">1. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Metal ions **<span style="font-family: Arial,sans-serif; font-size: 11pt;">- a metal ion can be displaced from a solution of its salt by a metal **above it** in the activity series e.g. zinc will react with copper sulfate solution and produce copper metal and zinc nitrate solution

<span style="font-family: Arial,sans-serif; font-size: 11pt;">5. **Fe3+** - this metal ion is unusual in that although it can be reduced to the Fe metallic state, it can <span style="font-family: Arial,sans-serif; font-size: 11pt;"> also be easily reduced to the Fe2+ state.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">6. **Permanganate ion, MnO4** - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> - this purple ion is reduced to colourless Mn2+ ion if the reaction is carried out in acidic solution. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> MnO4 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) + 8H+(//aq//) + 5e - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Mn2+(//aq//) + 4H2O(//l//) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> **N.B.** If the reaction is carried out in neutral or slightly basic conditions then the reduction <span style="font-family: Arial,sans-serif; font-size: 11pt;">product is a brown solid, MnO2.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">This reaction requires an acid catalyst. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cr2O72 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) + 14H+(//aq//) + 6e - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2Cr3+(//aq//) + 7H2O(//l//)
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">7. Dichromate ion, Cr2O72 ** - <span style="font-family: Arial,sans-serif; font-size: 11pt;">- this orange ion is reduced to green Cr3+.

=<span style="font-family: Arial,sans-serif; font-size: 11pt;">Common Reductants = <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Zn(//s//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Zn2+(//aq//) + 2e -
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">1. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Metals **<span style="font-family: Arial,sans-serif; font-size: 11pt;">, especially those high on the activity series, are oxidised to metal ions.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">green, Fe2+, to orange, Fe3+.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">2. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Iron(II) ion, Fe2+, **<span style="font-family: Arial,sans-serif; font-size: 11pt;">can be easily oxidised to Fe3+. The colour change observed is from pale

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> oxidation of foods e.g. in dried fruits and wines.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">3. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Sulfur dioxide, SO2, and sulfite ion, SO32 ** - <span style="font-family: Arial,sans-serif; font-size: 11pt;">are both oxidised to sulfate ion, SO42 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">. All of these species are colourless. Sulfur dioxide is commonly used as a preservative as it delays

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> solution of Br2.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">4. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Halide ions **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> - Iodide ions, I - <span style="font-family: Arial,sans-serif; font-size: 11pt;">, are readily oxidised to I2. The colour change is from colourless to yellow/brown in aqueous solution. Similarly, colourless Br - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> ions are oxidised to an orange


 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">5. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hydrogen gas, H2 **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> which can be oxidised to water e.g. CuO(//s//) + H2(//g//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cu(//s//) + H2O(//g//)

<span style="font-family: Arial,sans-serif; font-size: 11pt;">6. **Carbon, C or carbon monoxide, CO**. Carbon is usually a black solid while carbon monoxide is a colourless, poisonous gas. These are commonly used to reduce metal ores to obtain the metal. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe2O3(//s//) + 3CO(//g//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2Fe(//s//) + 3CO2(//g//)

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Electrolysis <span style="font-family: Arial,sans-serif; font-size: 11pt;">Electrolysis refers to a redox reaction that is made to occur by passing a direct electric current through an **electrolyte**, a liquid that conducts electricity. An electrolyte contains ions that are free to move and is either a molten ionic compound or an aqueous solution of an ionic salt.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Electrodes are immersed in the electrolyte and connected to the DC power supply (battery). The negative electrode is called the **cathode** and the positive electrode is called the **anode**. The electrodes are often inert e.g. graphite or platinum, and are not involved in the electrolysis reaction itself.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The whole set-up of electrolyte and electrodes is called an electrolytic cell. When it is working:
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;"> cations **<span style="font-family: Arial,sans-serif; font-size: 11pt;">move towards the cathode and are **reduced**
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;"> anions **<span style="font-family: Arial,sans-serif; font-size: 11pt;">migrate to the anode and are **oxidised.**

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Reaction only occurs at the surface of the electrodes. <span style="font-family: Arial,sans-serif; font-size: 11pt;">In the electrolytic cell shown <span style="font-family: Arial,sans-serif; font-size: 11pt;"> **D.C.**

**<span style="font-family: Arial,sans-serif; font-size: 11pt;"> cathode anode **
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 * ||^  ||   || [[image:https://skcchemistry.wikispaces.com/site/embedthumbnail/placeholder?w=200&h=50 width="200" height="50"]] ||
 * ||^  ||   || [[image:https://skcchemistry.wikispaces.com/site/embedthumbnail/placeholder?w=200&h=50 width="200" height="50"]] ||


 * <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Na+ Cl ** -

<span style="font-family: Arial,sans-serif; font-size: 11pt;">cathode reaction: Na+(//l//) + 1e - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Na(//s//) anode reaction: 2Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//l//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cl2(//g//) + 2e -

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Overall balanced equation for the cell reaction: <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//l//) + 2Na+(//l//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cl2(//g//) + 2Na(//s//) <span style="font-family: Arial,sans-serif; font-size: 11pt;">In the electrolysis shown above the salt was molten. If the electrolyte is an aqueous solution then it is possible that the water will undergo oxidation (to produce oxygen gas) or reduction (to produce hydrogen gas). It depends on the ease with which the ions of the salt are electrolysed compared to the water, as well as their concentration e.g. Electrolysis of a concentrated aqueous solution of NaCl (called brine) results in oxidation of Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> to chlorine gas, Cl2, at the anode but reduction of water to hydrogen gas at the cathode. This is because Na+ is not easily reduced - it is above H in the activity series.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">If the solution of NaCl is very dilute then oxidation of water occurs at the anode producing O2 gas rather than Cl2 gas.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">In contrast, electrolysis of aqueous copper chloride produces Cu(s) and Cl2(g), since Cu is below H in the activity series. Electrolysis of aqueous copper sulfate produces pink copper metal on the cathode but oxygen gas at the anode since water is more easily oxidised than the sulfate ion. (Note: In SO42 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> the S is already in its maximum oxidation state of +6 so it cannot be oxidised further.)

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Electrolysis is an important industrial process used to extract elements. Examples are: <span style="font-family: Arial,sans-serif; font-size: 11pt;">1) The electrolytic extraction of Al from alumina, Al2O3, originally found in the ore bauxite which is mined in Northern Queensland. The electrolysis occurs at Tiwai Point, Southland. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Reactions occurring during electrolysis are: <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cathode: Al3+ + 3e - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Al (reduction) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Anode: 2O2 - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> O2 + 4e - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> (oxidation)

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The carbon anodes that are used (as metal electrodes would alloy with the molten aluminium that is produced) have to be replaced periodically because they react with the oxygen produced to form carbon dioxide.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">2) The electrolytic extraction of magnesium from sea water. <span style="font-family: Arial,sans-serif; font-size: 11pt;">The seawater is first concentrated by evaporation during which most of the sodium chloride present precipitates out. Solid magnesium chloride is obtained by precipitation with hydroxide followed by reaction with hydrochloric acid. Reactions occurring during electrolysis are:

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cathode: Mg2+ + 2e - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Mg (reduction) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Anode: 2Cl - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cl2 + 2e - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> (oxidation)