Oxidation+Reduction+Reactions

An oxidation-reduction reaction (or redox reaction) is one that involves the transfer of electrons from one species to another. Oxidation was originally defined as a reaction with oxygen. For example Mg reacts with O2 to form magnesium oxide, MgO. 2Mg(//s//) + O2(//g//) ®  2 MgO(//s//) Similarly reduction was the removal of oxygen eg. CO reduces Fe2O3 to produce Fe and CO2. Fe2O3(//s//) + 3CO(//g//) ®  2Fe(//s//) + 3CO2(//g//) There are also, however, reactions that are classified as redox reactions that do not involve oxygen eg. the displacement by zinc metal of copper metal from a solution of copper sulfae. 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 ** - ** ®  Cu 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;">Redox Definitions **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Rather than consider gain or loss of oxygen or electrons, a better method is to use a system of **oxidation numbers (states).** In all redox reactions there is a change in oxidation numbers.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">An increase in oxidation number corresponds to oxidation
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">A decrease in oxidation number corresponds to reduction.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The oxidation number (or state) is a number that can be assigned to each individual atom in an element compound or ion, using a set of rules. These rules are as follows: <span style="font-family: Arial,sans-serif; font-size: 11pt;">eg. In Na+ the oxidation number is +1, in O**2** ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> the oxidation number is -2. <span style="font-family: Arial,sans-serif; font-size: 11pt;">In an ionic compound containing monatomic ions the ions have the same oxidation numbers as they would alone eg. in Na2O the oxidation numbers are still +1 for Na+ and -2 for O**2** ** - ****<span style="font-family: Arial,sans-serif; font-size: 11pt;">. ** <span style="font-family: Arial,sans-serif; font-size: 11pt;">4. In compounds each oxygen atom has oxidation number of -2 (except in peroxides eg. H2O2 when it is -1). <span style="font-family: Arial,sans-serif; font-size: 11pt;">5. 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;">Example: <span style="font-family: Arial,sans-serif; font-size: 11pt;">a) Find the oxidation number of S in H2SO4. <span style="font-family: Arial,sans-serif; font-size: 11pt;">(2 x +1) + (1 x ?) + (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;">b) 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 = = +6 <span style="font-family: Arial,sans-serif; font-size: 11pt;">c) In an ionic compound such as NH4NO3 the oxidation numbers of N are determined by first separating into ions NH4+ and NO3- and then finding the oxidation numbers. The values obtained are –3 in NH4+ and +5 in NO3 ** - ** //<span style="font-family: Arial,sans-serif; font-size: 11pt;">Exercise //<span style="font-family: Arial,sans-serif; font-size: 11pt;">: //<span style="font-family: Arial,sans-serif; font-size: 11pt;">1) Find the oxidation numbers of the S atoms in each of the following substances. // //<span style="font-family: Arial,sans-serif; font-size: 11pt;">SO2, H2SO4, SO3, S8, H2SO3, Na2S2O3 //
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Oxidation numbers **
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">1. The oxidation number of an atom in an element is **zero** eg. in H2 the oxidation number of H is 0.
 * 2) <span style="font-family: Arial,sans-serif; font-size: 11pt;">2. The oxidation number of an atom in a monatomic ion is the same as the charge on the ion.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">3. In compounds each hydrogen atom usually has an oxidation number of +1 (the exception is in the metal hydrides e.g.NaH where the oxidation number of H= -1).
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">6. In **polyatomic ions** the sum of the oxidation numbers of all the atoms is equal to the overall charge on the ion.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">2) Calculate the oxidation number of N in each of the following ions: //<span style="font-family: Arial,sans-serif; font-size: 11pt;">NO3 //** - **//<span style="font-family: Arial,sans-serif; font-size: 11pt;">, NH4+, NO2 //** - **//<span style="font-family: Arial,sans-serif; font-size: 11pt;">, N //**<span style="font-family: Arial,sans-serif; font-size: 11pt;">3 **** - **

//<span style="font-family: Arial,sans-serif; font-size: 11pt;">3) // //<span style="font-family: Arial,sans-serif; font-size: 11pt;">By assigning oxidation numbers show which of the following reactions is **not** a redox reaction. //

//<span style="font-family: Arial,sans-serif; font-size: 11pt;">(a) CuCO3 // ® //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> CuO + CO2 //

//<span style="font-family: Arial,sans-serif; font-size: 11pt;">(b) Cu + 2AgNO3 // ® //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cu(NO3)2 + 2Ag //

//<span style="font-family: Arial,sans-serif; font-size: 11pt;">(c) Cr2O //**<span style="font-family: Arial,sans-serif; font-size: 11pt;">7 ****<span style="font-family: Arial,sans-serif; font-size: 11pt;">2 **** - **//<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + 6Fe2+ // ® //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> 6Fe3+ + 2Cr3+ + 7H2O //

//<span style="font-family: Arial,sans-serif; font-size: 11pt;">(d) Cr2O7 //**<span style="font-family: Arial,sans-serif; font-size: 11pt;">2 **** - **//<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + 2OH //** - ** ® //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2CrO //**<span style="font-family: Arial,sans-serif; font-size: 11pt;">4 ****<span style="font-family: Arial,sans-serif; font-size: 11pt;">2 **** - **//<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + H2O //

<span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTE: For the elements of the //p// block (Groups 13 – 18) the maximum oxidation number possible is the effective charge when the atom effectively loses all its valence electrons (by bonding with a more electronegative atom). The minimum value is when the atom gains electrons to complete its outer shell. For example, sulfur is in Group 16 and therefore has a minimum value of –2 (as in H2S) and a maximum velue of +6 (as in H2SO4). This is why H2S can only be oxidised while SO42 - <span style="font-family: Arial,sans-serif; font-size: 11pt;"> can only be reduced while SO2 (oxidation number +4) can be either oxidised or reduced.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Balancing Redox Equations <span style="font-family: Arial,sans-serif; font-size: 11pt;">The following method for balancing redox equations is commonly called the ion-electron half-equation method. You should use this method if you cannot balance both the elements **and** the charge by inspection or by trial and error.


 * <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. In the reduction reaction the species reduced will be the reactant that has an oxidation number that decreases. The opposite is true for the oxidation half reaction.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">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;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ and Cr2O72 - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cr3+ <span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ and Cr2O72 - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> **2**Cr3+ <span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2 ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ and Cr2O72 - ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> **2**Cr3+ + **7H2O** <span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Fe3+ and 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="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;"> and 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;">**OIL RIG** <span style="font-family: Arial,sans-serif; font-size: 11pt;">**o**xidation **i**s **l**oss **r**eduction **i**s **g**ain <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;"> 6 Fe3+ + 6**e** ** - ** <span style="font-family: Arial,sans-serif; font-size: 11pt;">Added to Cr2O7**2** ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + 14H+ + 6e ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2Cr3+ + 7H2O gives the final equation <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;">1. MnO4 // - //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + SO2 // ® //<span style="font-family: Arial,sans-serif; font-size: 11pt;"> Mn2+ + SO4 //**<span style="font-family: Arial,sans-serif; font-size: 11pt;">2 **** - **
 * <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;">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;">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;">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;">Step 6 **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> - To obtain an overall balanced equation the two half equations must be added together. Before doing this the equations may have to be multiplied so that the number of electrons in each half-equation is the same. In this way, the electrons will be removed from the final equation.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">6Fe2+ + **<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;"> **2Cr3+ + 7H2O + 6 Fe3+**
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Exercises **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> - Balance each of the following equations.

//<span style="font-family: Arial,sans-serif; font-size: 11pt;">2. S2O3 //**<span style="font-family: Arial,sans-serif; font-size: 11pt;">2 **** - **//<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + I2 // ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> //I// ** - **//<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + S4O6 //**<span style="font-family: Arial,sans-serif; font-size: 11pt;">2 **** - **

<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 ** - ** <span style="font-family: Arial,sans-serif; font-size: 11pt;">Both Zn2+ and Mg2+ are colourless while Fe2+ is pale green in solution.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Common Reductants **
 * 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. Common examples are Zn (forms Zn2+), Mg (forms Mg2+) and Fe (forms Fe2+)


 * 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 green Fe2+ to orange Fe3+. This occurs when a precipitate of dark green Fe(OH)2 is exposed to air.

<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 oxidation of foods such as dried fruits and wines. Solutions of SO2 in water form sulfurous acid, H2SO3 which is a weak acid and dissociates to form HSO3 ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> and SO3**2** ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> ions.
 * 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, SO4**2** ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">.


 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">4. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Iodide ions, I **** - ****<span style="font-family: Arial,sans-serif; font-size: 11pt;"> and bromide ions, Br **** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">are both colourless and are readily oxidised to the halogens I2 and Br2 respectively. When I2 is formed the solution goes a brown/yellow colour in aqueous solution. Because of its limited solubility solid iodine may form as a black precipitate of I2. The formation of Br2 results in the solution turning an orange colour.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">5**. Thiosulfate ions, S2O32** ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">- This colourless ion is commonly used in the redox titration which determines the concentration of iodine in a solution. The thiosulfate ion is oxidised to S4O6**2** ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> (tetrathionate ion). <span style="font-family: Arial,sans-serif; font-size: 11pt;">I2 + S2O3**2** ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> S4O6**2** ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + 2I ** - ** <span style="font-family: Arial,sans-serif; font-size: 11pt;">6. **Carbon, C, and carbon monoxide, CO -** Carbon is commonly used as a reductant in the form of graphite or coal (both black solids). In limited air it is oxidised to CO (a colourless, poisonous gas), and in a plentiful supply of air it is oxidised to CO2 (also a colourless gas). <span style="font-family: Arial,sans-serif; font-size: 11pt;">CO (made from combustion of coal) is used as the reductant at the Glenbrook Steel Mill where iron(III) oxide, Fe2O3 (from iron sand) is reduced to Fe and CO is oxidised to CO2. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Fe2O3 + 3CO ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2Fe + 3CO2
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">7. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hydrogen gas, H2 **<span style="font-family: Arial,sans-serif; font-size: 11pt;">- This colourless gas can be used as a reductant, commonly combining with oxygen to form water, H2O.


 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">8. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Oxalic acid, H2C2O4 **<span style="font-family: Arial,sans-serif; font-size: 11pt;">– This colourless, posonous substance is a weak acid that is oxidised to carbon dioxide CO2. Its anion is the oxalate ion C2O4**2** - <span style="font-family: Arial,sans-serif; font-size: 11pt;">, found in salts such as sodium oxalate.


 * 1) **<span style="font-family: Arial,sans-serif; font-size: 11pt;">9. ** **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hydrogen sulfide, H2S **<span style="font-family: Arial,sans-serif; font-size: 11pt;">– This gas can be used as a reductant as it can be oxidised to sulfur species in a higher oxidation state eg SO2.


 * 1) **<span style="font-family: Arial,sans-serif; font-size: 11pt;">10. ** **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hydrogen peroxide, H2O2 **<span style="font-family: Arial,sans-serif; font-size: 11pt;">– When acting as a reductant it is oxidised to O2. This may be seen as bubbles of gas. Note that H2O2 can also act as an oxidant (see below).

=<span style="font-family: Arial,sans-serif; font-size: 11pt;">Common Oxidants =
 * 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;"> is involved in all burning reactions producing the oxide ion, O**2** ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">. Both species are colourless so no colour change is observed.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">potassium K Kate’s <span style="font-family: Arial,sans-serif; font-size: 11pt;">sodium Na Nanny <span style="font-family: Arial,sans-serif; font-size: 11pt;">calcium Ca Can. <span style="font-family: Arial,sans-serif; font-size: 11pt;">magnesium Mg Manage <span style="font-family: Arial,sans-serif; font-size: 11pt;">aluminium Al A <span style="font-family: Arial,sans-serif; font-size: 11pt;">zinc Zn Zebra <span style="font-family: Arial,sans-serif; font-size: 11pt;">iron Fe For <span style="font-family: Arial,sans-serif; font-size: 11pt;">tin Sn She <span style="font-family: Arial,sans-serif; font-size: 11pt;">lead Pb Punishes <span style="font-family: Arial,sans-serif; font-size: 11pt;">copper Cu Cruelly <span style="font-family: Arial,sans-serif; font-size: 11pt;">mercury Hg How <span style="font-family: Arial,sans-serif; font-size: 11pt;">silver Ag Agonisingly <span style="font-family: Arial,sans-serif; font-size: 11pt;">gold Au Awful
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">2. **<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. Metals above hydrogen in the activity series will react with H+ in dilute acid or water. Generally the more active the metal the more violent the reaction.
 * || <span style="font-family: Arial,sans-serif; font-size: 11pt;">Most reactive metals react violently with water. They are not found naturally. ||  ||
 * //<span style="font-family: Arial,sans-serif; font-size: 11pt;">Activity Series: //**
 * || <span style="font-family: Arial,sans-serif; font-size: 11pt;">React slowly with water and acid. ||  ||
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Hydrogen H Him **
 * || <span style="font-family: Arial,sans-serif; font-size: 11pt;">Metals below H are unreactive. ||  ||

<span style="font-family: Arial,sans-serif; font-size: 11pt;">MnO4 ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + 8H+ + 5e ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Mn2+ + 4H2O <span style="font-family: Arial,sans-serif; font-size: 11pt;">MnO2 + 4H+ + 2e ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Mn2+ + 2H2O <span style="font-family: Arial,sans-serif; font-size: 11pt;">If the reaction of MnO4- is carried out in neutral or slightly basic conditions then the reduction product is a brown solid, MnO2. <span style="font-family: Arial,sans-serif; font-size: 11pt;">MnO4 ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> //+// 2H2O + 3e ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> MnO2 + 4OH ** - ** <span style="font-family: Arial,sans-serif; font-size: 11pt;">In strongly basic conditions the permanganate ion is reduced to the green manganate ion <span style="font-family: Arial,sans-serif; font-size: 11pt;">MnO4 ** - **//<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + //<span style="font-family: Arial,sans-serif; font-size: 11pt;">e ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> MnO4**2** ** - ** <span style="font-family: Arial,sans-serif; font-size: 11pt;">Cr2O7**2** ** - ** <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="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;">7. **Metal ions (Fe3+ and Cu2+)** - A metal ion can be displaced from a solution of its salt by a metal **above it** in the activity series. The orange Fe3+ ion usually undergoes reduction to the pale green Fe2+ ion.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">3. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Halogens **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> - 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;">. Because of its oxidising properties Cl2 is used as a disinfectant and to sterilise swimming pools.
 * 2) <span style="font-family: Arial,sans-serif; font-size: 11pt;">4. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Permanganate ion, MnO4- **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> **and** **manganese dioxide MnO2** - The purple ion MnO4- and the brown solid MnO2 are both reduced to colourless Mn2+ ion if the reaction is carried out in acidic solution.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">5. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Dichromate ion, Cr2O72 **** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">- This orange ion is reduced to green Cr3+. The reaction requires an acid catalyst.
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">6. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hydrogen peroxide, H2O2 **<span style="font-family: Arial,sans-serif; font-size: 11pt;">- When colourless H2O2 acts as an oxidant it is reduced to water.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Blue Cu2+ may be reduced to either Cu+ or pink Cu metal depending on the reducing agent used. The reduction of Cu2+ in the redox reaction of aldehydes with Benedict’s solution produces the red precipitate of Cu2O (copper I oxide).

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The reduction of Cu2+ using I- produces a white precipitate of CuI (copper I iodide). <span style="font-family: Arial,sans-serif; font-size: 11pt;">Cu2+ + CH3CHO ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cu2O + CH3CO2H <span style="font-family: Arial,sans-serif; font-size: 11pt;">Cu2+ + I ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> CuI + I2 <span style="font-family: Arial,sans-serif; font-size: 11pt;">8. **Concentrated nitric acid, HNO3** - When colourless conc nitric acid is used as an oxidant. When reacted with copper metal, the observed product is brown nitrogen dioxide gas, NO2. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Cu + 2HNO3 + 2H+ ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2NO2 + Cu2+ + 2H2O <span style="font-family: Arial,sans-serif; font-size: 11pt;">More dilute nitric acid solutions may produce colourless NO gas but this is rapidly oxidised to brown NO2 by oxygen in the air. <span style="font-family: Arial,sans-serif; font-size: 11pt;">2IO3 ** - ** <span style="font-family: Arial,sans-serif; font-size: 11pt;">+ 12H+ + 10e ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> I2 + 6H2O <span style="font-family: Arial,sans-serif; font-size: 11pt;">2OCl ** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> + 4H+ + 2e ** - ** ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Cl2 + 2H2O
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">9. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Iodate ion, IO3 **** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">- This ion is colourless and can be reduced to form the halogen, I2 (a black solid or when dissolved in water a brown aqueous solution).
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">10. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Hypochlorite ion, OCl **** - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">- A colourless ion that can be reduced to form chlorine, Cl2, a yellow-green gas.

**<span style="font-family: Arial,sans-serif; font-size: 11pt;">Exercise **
//<span style="font-family: Arial,sans-serif; font-size: 11pt;">Compare the reaction of hydrogen peroxide with potassium permanganate under both neutral and acidic conditions. Write balanced equations for the reactions that occur, and note any observations that would be made. //