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## REDOX A guide for A level students

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**REDOX**A guide for A level students 2008 SPECIFICATIONS KNOCKHARDY PUBLISHING**KNOCKHARDY PUBLISHING**REDOX INTRODUCTION This Powerpoint show is one of several produced to help students understand selected topics at AS and A2 level Chemistry. It is based on the requirements of the AQA and OCR specifications but is suitable for other examination boards. Individual students may use the material at home for revision purposes or it may be used for classroom teaching if an interactive white board is available. Accompanying notes on this, and the full range of AS and A2 topics, are available from the KNOCKHARDY SCIENCE WEBSITE at... www.knockhardy.org.uk/sci.htm Navigation is achieved by... either clicking on the grey arrows at the foot of each page or using the left and right arrow keys on the keyboard**REDOX**• CONTENTS • Definitions of oxidation and reduction • Calculating oxidation state • Use of H, O and F in calculating oxidation state • Naming compounds • Redox reactions • Balancing ionic half equations • Combining half equations to form a redox equation • Revision check list**REDOX**• Before you start it would be helpful to… • Recall the layout of the periodic table • Be able to balance simple equations**OXIDATION & REDUCTION - Definitions**OXIDATION GAIN OF OXYGEN 2Mg + O2——> 2MgO magnesium has been oxidised as it has gained oxygen REMOVAL (LOSS) OF HYDROGEN C2H5OH ——> CH3CHO + H2 ethanol has been oxidised as it has ‘lost’ hydrogen**OXIDATION & REDUCTION - Definitions**REDUCTION GAIN OF HYDROGEN C2H4 + H2 ——> C2H6 ethene has been reduced as it has gained hydrogen REMOVAL (LOSS) OF OXYGEN CuO + H2 ——> Cu + H2O copper(II) oxide has been reduced as it has ‘lost’ oxygen However as chemistry became more sophisticated, it was realised that another definition was required**OXIDATION & REDUCTION - Definitions**OXIDATION AND REDUCTION IN TERMS OF ELECTRONS Oxidation and reduction are not only defined as changes in O and H ... OXIDATIONRemoval (loss) of electrons ‘OIL’ species will get less negative or more positive REDUCTIONGain of electrons ‘RIG’ species will become more negative or less positive REDOXWhen reduction and oxidation take place**OXIDATION & REDUCTION - Definitions**OXIDATION AND REDUCTION IN TERMS OF ELECTRONS Oxidation and reduction are not only defined as changes in O and H ... OXIDATIONRemoval (loss) of electrons ‘OIL’ species will get less negative or more positive REDUCTIONGain of electrons ‘RIG’ species will become more negative or less positive REDOXWhen reduction and oxidation take place OIL -Oxidation Is the Loss of electrons RIG -Reduction Is the Gain of electrons**OXIDATION STATES**Used to... tell if oxidation or reduction has taken place work out what has been oxidised and/or reduced construct half equations and balance redox equations ATOMS AND SIMPLE IONS The number of electrons which must be added or removed to become neutral atomsNa in Na = 0 neutral already ... no need to add any electrons cationsNa in Na+ = +1 need to add 1 electron to make Na+ neutral anionsCl in Cl¯ = -1 need to take 1 electron away to make Cl¯ neutral**OXIDATION STATES**Used to... tell if oxidation or reduction has taken place work out what has been oxidised and/or reduced construct half equations and balance redox equations ATOMS AND SIMPLE IONS The number of electrons which must be added or removed to become neutral atomsNa in Na = 0 neutral already ... no need to add any electrons cationsNa in Na+ = +1 need to add 1 electron to make Na+ neutral anionsCl in Cl¯ = -1 need to take 1 electron away to make Cl¯ neutral Q. What are the oxidation states of the elements in the following? a) C b) Fe3+ c) Fe2+ d) O2- e) He f) Al3+**OXIDATION STATES**Used to... tell if oxidation or reduction has taken place work out what has been oxidised and/or reduced construct half equations and balance redox equations ATOMS AND SIMPLE IONS The number of electrons which must be added or removed to become neutral atomsNa in Na = 0 neutral already ... no need to add any electrons cationsNa in Na+ = +1 need to add 1 electron to make Na+ neutral anionsCl in Cl¯ = -1 need to take 1 electron away to make Cl¯ neutral Q. What are the oxidation states of the elements in the following? a) C (0) b) Fe3+(+3) c) Fe2+ (+2) d) O2-(-2) e) He (0) f) Al3+ (+3)**OXIDATION STATES**MOLECULES The SUM of the oxidation states adds up to ZERO ELEMENTSH in H2 = 0 both are the same and must add up to Zero COMPOUNDSC in CO2 = +4 O in CO2 = -2 1 x +4 and 2 x -2 = Zero**OXIDATION STATES**MOLECULES The SUM of the oxidation states adds up to ZERO ELEMENTSH in H2 = 0 both are the same and must add up to Zero COMPOUNDSC in CO2 = +4 O in CO2 = -2 1 x +4 and 2 x -2 = Zero Explanation • because CO2 is a neutralmolecule, the sumoftheoxidationstatesmust be zero • for this, oneelementmusthave a positiveOS and the othermustbenegative**OXIDATION STATES**MOLECULES The SUM of the oxidation states adds up to ZERO ELEMENTS H in H2 = 0 both are the same and must add up to Zero COMPOUNDS C in CO2 = +4 O in CO2 = -2 1 x +4 and 2 x +2 = Zero • HOW DO YOU DETERMINE WHICH IS THE POSITIVE ONE? • the more electronegative species will have the negativevalue • electronegativity increases across a period and decreases down a group • O is further to the right than C in the periodic table so it has the negative value**OXIDATION STATES**MOLECULES The SUM of the oxidation states adds up to ZERO ELEMENTS H in H2 = 0 both are the same and must add up to Zero COMPOUNDS C in CO2 = +4 O in CO2 = -2 1 x +4 and 2 x +2 = Zero • HOW DO YOU DETERMINE THE VALUE OF • AN ELEMENT’S OXIDATION STATE? • from its positionintheperiodictable and/or • the other element(s) present in the formula**OXIDATION STATES**COMPLEX IONS The SUM of the oxidation states adds up to THE CHARGE e.g. NO3-sum of the oxidation states = - 1 SO42- sum of the oxidation states = - 2 NH4+ sum of the oxidation states = +1 Examples in SO42- the oxidation state of S = +6 there is ONE S O = -2 there are FOUR O’s +6 + 4(-2) = -2 so the ion has a 2- charge**OXIDATION STATES**COMPLEX IONS The SUM of the oxidation states adds up to THE CHARGE e.g. NO3-sum of the oxidation states = - 1 SO42- sum of the oxidation states = - 2 NH4+ sum of the oxidation states = +1 Examples • What is the oxidation state (OS) of Mn in MnO4¯ ? • the oxidation state of oxygen in most compounds is - 2 • there are 4 O’s so the sum of its oxidation states - 8 • overall charge on the ion is - 1 • therefore the sum of all the oxidation states must add up to - 1 • the oxidation states of Mn four O’s must therefore equal - 1 • therefore the oxidation state of Mn in MnO4¯is +7 • +7 + 4(-2) = - 1**OXIDATION STATES**CALCULATING OXIDATION STATE - 1 Many elements can exist in more than one oxidation state In compounds, certain elements are used as benchmarks to work out other values HYDROGEN+1 except0 atom (H) and molecule (H2) -1 hydride ion, H¯ in sodium hydride NaH OXYGEN -2 except 0 atom (O) and molecule (O2) -1 in hydrogen peroxide, H2O2 +2 in F2O FLUORINE -1 except 0 atom (F) and molecule (F2)**OXIDATION STATES**CALCULATING OXIDATION STATE - 1 Many elements can exist in more than one oxidation state In compounds, certain elements are used as benchmarks to work out other values HYDROGEN+1 except0 atom (H) and molecule (H2) -1 hydride ion, H¯ in sodium hydride NaH OXYGEN -2 except 0 atom (O) and molecule (O2) -1 in hydrogen peroxide, H2O2 +2 in F2O FLUORINE -1 except 0 atom (F) and molecule (F2) Q.Give the oxidation state of the element other than O, H or F in... SO2NH3NO2NH4+IF7Cl2O7 NO3¯ NO2¯ SO32-S2O32-S4O62-MnO42- What is odd about the value of the oxidation state of S in S4O62- ?**OXIDATION STATES**A. The oxidation states of the elements other than O, H or F are SO2 O = -2 2 x -2 = - 4 overall neutral S = +4 NH3 H = +1 3 x +1 = +3 overall neutral N = - 3 NO2 O = -22 x -2 = - 4 overall neutral N = +4 NH4+ H = +1 4 x +1 = +4 overall +1 N = - 3 IF7 F = -1 7 x -1 = - 7 overall neutral I = +7 Cl2O7 O = -27 x -2 = -14 overall neutral Cl = +7 (14/2) NO3¯ O = -23 x -2 = - 6 overall -1 N = +5 NO2¯ O = -22 x -2 = - 4 overall -1 N = +3 SO32- O = -23 x -2 = - 6 overall -2 S = +4 S2O32- O = -23 x -2 = - 6 overall -2 S = +2 (4/2) S4O62- O = -26 x -2 = -12 overall -2 S = +2½ ! (10/4) MnO42- O = -24 x -2 = - 8 overall -2 Mn = +6 What is odd about the value of the oxidation state of S in S4O62- ? An oxidation state must be a whole number (+2½ is the average value)**OXIDATION STATES**CALCULATING OXIDATION STATE - 2 The position of an element in the periodic table can act as a guide METALS• have positive values in compounds • value is usually that of the Group Number Al is+3 • where there are several possibilities the values go no higher than the Group No. Sn can be+2 or +4 Mn can be +2,+4,+6,+7 NON-METALS • mostly negative based on their usual ion Cl usually-1 • can have values up to their Group No. Cl +1 +3 +5 or +7**OXIDATION STATES**CALCULATING OXIDATION STATE - 2 The position of an element in the periodic table can act as a guide METALS• have positive values in compounds • value is usually that of the Group Number Al is+3 • where there are several possibilities the values go no higher than the Group No. Sn can be+2 or +4 Mn can be +2,+4,+6,+7 NON-METALS • mostly negative based on their usual ion Cl usually-1 • can have values up to their Group No. Cl +1 +3 +5or+7 Q. What is the theoretical maximum oxidation state of the following elements? Na P Ba Pb S Mn Cr What will be the usual and the maximum oxidation state in compounds of? Li Br Sr O B N+1**OXIDATION STATES**CALCULATING OXIDATION STATE - 2 The position of an element in the periodic table can act as a guide A.What is the theoretical maximum oxidation state of the following elements? Na P Ba Pb S Mn Cr +1 +5 +2 +4 +6 +7 +6 What will be the usual and the maximum oxidation state in compounds of? Li Br Sr O B N USUAL+1 -1 +2 -2 +3 -3 or +5 MAXIMUM+1 +7 +2 +6 +3 +5**OXIDATION STATES**CALCULATING OXIDATION STATE - 2 Q. What is the oxidation state of each element in the following compounds/ions ? CH4 PCl3 NCl3 CS2 ICl5 BrF3 PCl4+ H3PO4 NH4Cl H2SO4 MgCO3 SOCl2**OXIDATION STATES**CALCULATING OXIDATION STATE - 2 Q. What is the oxidation state of each element in the following compounds/ions ? CH4C = - 4 H = +1 PCl3P = +3 Cl = -1 NCl3N = +3 Cl = -1 CS2C = +4 S = -2 ICl5I = +5 Cl = -1 BrF3Br = +3 F = -1 PCl4+P = +4 Cl = -1 H3PO4P = +5 H = +1 O = -2 NH4Cl N = -3 H = +1 Cl = -1 H2SO4S = +6 H = +1 O = -2 MgCO3Mg = +2 C = +4 O = -2 SOCl2S = +4 Cl = -1 O = -2**OXIDATION STATES**THE ROLE OF OXIDATION STATE IN NAMING SPECIES To avoid ambiguity, the oxidation state is often included in the name of a species manganese(IV) oxide shows that Mn is in the +4 oxidation state in MnO2 sulphur(VI) oxide for SO3 S is in the +6 oxidation state dichromate(VI) for Cr2O72- Cr is in the +6 oxidation state phosphorus(V) chloride for PCl5 P is in the +5 oxidation state phosphorus(III) chloride for PCl3 P is in the +3 oxidation state Q. Name the following... PbO2 SnCl2 SbCl3 TiCl4 BrF5**OXIDATION STATES**THE ROLE OF OXIDATION STATE IN NAMING SPECIES To avoid ambiguity, the oxidation state is often included in the name of a species manganese(IV) oxide shows that Mn is in the +4 oxidation state in MnO2 sulphur(VI) oxide for SO3 S is in the +6 oxidation state dichromate(VI) for Cr2O72- Cr is in the +6 oxidation state phosphorus(V) chloride for PCl5 P is in the +5 oxidation state phosphorus(III) chloride for PCl3 P is in the +3 oxidation state Q. Name the following... PbO2lead(IV) oxide SnCl2tin(II) chloride SbCl3antimony(III) chloride TiCl4titanium(IV) chloride BrF5bromine(V) fluoride**REDOX REACTIONS**OXIDATION AND REDUCTION IN TERMS OF ELECTRONS Oxidation and reduction are not only defined as changes in O and H REDOXWhen reduction and oxidation take place OXIDATIONRemoval (loss) of electrons ‘OIL’ species will get less negative or more positive REDUCTIONGain of electrons ‘RIG’ species will become more negative or less positive**REDOX REACTIONS**OXIDATION AND REDUCTION IN TERMS OF ELECTRONS Oxidation and reduction are not only defined as changes in O and H REDOXWhen reduction and oxidation take place OXIDATIONRemoval (loss) of electrons ‘OIL’ species will get less negative or more positive REDUCTIONGain of electrons ‘RIG’ species will become more negative or less positive REDUCTION in O.S. Species has been REDUCED e.g. Cl is reduced to Cl¯ (0 to -1) INCREASE in O.S.Species has been OXIDISED e.g. Na is oxidised to Na+ (0 to +1)**REDOX REACTIONS**OXIDATION AND REDUCTION IN TERMS OF ELECTRONS REDUCTION in O.S. INCREASE in O.S. Species has been REDUCEDSpecies has been OXIDISED Q. State if the changes involve oxidation (O) or reduction (R) or neither (N) Fe2+ —> Fe3+ I2 —> I¯ F2 —> F2O C2O42- —> CO2 H2O2 —> O2 H2O2 —> H2O Cr2O72- —> Cr3+ Cr2O72- —> CrO42- SO42- —> SO2**REDOX REACTIONS**OXIDATION AND REDUCTION IN TERMS OF ELECTRONS REDUCTION in O.S. INCREASE in O.S. Species has been REDUCEDSpecies has been OXIDISED Q. State if the changes involve oxidation (O) or reduction (R) or neither (N) Fe2+ —> Fe3+ O +2 to +3 I2 —> I¯ R 0 to -1 F2 —> F2O R 0 to -1 C2O42- —> CO2 O +3 to +4 H2O2 —> O2 O -1 to 0 H2O2 —> H2O R -1 to -2 Cr2O72- —> Cr3+ R +6 to +3 Cr2O72- —> CrO42- N +6 to +6 SO42- —> SO2 R +6 to +4**BALANCING REDOX HALF EQUATIONS**1 Work out formulae of the species before and after the change; balance if required 2 Work out oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If equation still doesn’t balance, add sufficient water molecules to one side**BALANCING REDOX HALF EQUATIONS**1 Work out formulae of the species before and after the change; balance if required 2 Work out oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If equation still doesn’t balance, add sufficient water molecules to one side Example 1 Iron(II) being oxidised to iron(III) Step 1 Fe2+ ——> Fe3+ Step 2 +2+3 Step 3 Fe2+ ——> Fe3+ + e¯ now balanced An electron (charge -1) is added to the RHS of the equation... this balances the oxidation state change i.e. (+2) ——> (+3) + (-1) As everything balances, there is no need to proceed to Steps 4 and 5**BALANCING REDOX HALF EQUATIONS**1 Work out formulae of the species before and after the change; balance if required 2 Work out oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If equation still doesn’t balance, add sufficient water molecules to one side Example 2 MnO4¯ being reduced to Mn2+ in acidic solution**BALANCING REDOX HALF EQUATIONS**1 Work out formulae of the species before and after the change; balance if required 2 Work out oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If equation still doesn’t balance, add sufficient water molecules to one side Example 2 MnO4¯ being reduced to Mn2+ in acidic solution Step 1 MnO4¯ ———> Mn2+ No need to balance Mn; equal numbers**BALANCING REDOX HALF EQUATIONS**1 Work out formulae of the species before and after the change; balance if required 2 Work out oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If equation still doesn’t balance, add sufficient water molecules to one side Example 2 MnO4¯ being reduced to Mn2+ in acidic solution Step 1 MnO4¯ ———> Mn2+ Step 2 +7 +2 Overall charge on MnO4¯ is -1; sum of the OS’s of all atoms must add up to -1 Oxygen is in its usual oxidation state of -2; four oxygen atoms add up to -8 To make the overall charge -1, Mn must be in oxidation state +7 ... [+7 + (4x -2) = -1]**BALANCING REDOX HALF EQUATIONS**1 Work out formulae of the species before and after the change; balance if required 2 Work out oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If equation still doesn’t balance, add sufficient water molecules to one side Example 2 MnO4¯ being reduced to Mn2+ in acidic solution Step 1 MnO4¯ ———> Mn2+ Step 2 +7 +2 Step 3 MnO4¯ + 5e¯ ———> Mn2+ The oxidation states on either side are different; +7 —> +2 (REDUCTION) To balance; add 5 negative charges to the LHS [+7 + (5 x -1) = +2] You must ADD 5 ELECTRONS to the LHS of the equation**BALANCING REDOX HALF EQUATIONS**1 Work out formulae of the species before and after the change; balance if required 2 Work out oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If equation still doesn’t balance, add sufficient water molecules to one side Example 2 MnO4¯ being reduced to Mn2+ in acidic solution Step 1 MnO4¯ ———> Mn2+ Step 2 +7 +2 Step 3 MnO4¯ + 5e¯ ———> Mn2+ Step 4 MnO4¯ + 5e¯ + 8H+ ———> Mn2+ Total charges on either side are not equal; LHS = 1- and 5- = 6- RHS = 2+ Balance them by adding 8 positive charges to the LHS [ 6- + (8 x 1+) = 2+ ] You must ADD 8 PROTONS (H+ ions) to the LHS of the equation**BALANCING REDOX HALF EQUATIONS**1 Work out formulae of the species before and after the change; balance if required 2 Work out oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If equation still doesn’t balance, add sufficient water molecules to one side Example 2 MnO4¯ being reduced to Mn2+ in acidic solution Step 1 MnO4¯ ———> Mn2+ Step 2 +7 +2 Step 3 MnO4¯ + 5e¯ ———> Mn2+ Step 4 MnO4¯ + 5e¯ + 8H+ ———> Mn2+ Step 5 MnO4¯ + 5e¯ + 8H+ ———> Mn2+ + 4H2O now balanced Everything balances apart from oxygen and hydrogen O LHS = 4 RHS = 0 H LHS = 8 RHS = 0 You must ADD 4 WATER MOLECULES to the RHS; the equation is now balanced**BALANCING REDOX HALF EQUATIONS**Watch out for cases when the species is present in different amounts on either side of the equation ... IT MUST BE BALANCED FIRST Example 3 Cr2O72- being reduced to Cr3+ in acidic solution Step 1 Cr2O72- ———> Cr3+ there are two Cr’s on LHS Cr2O72- ———> 2Cr3+ both sides now have 2 Step 2 2 @ +6 2 @ +3 both Cr’s are reduced Step 3 Cr2O72- + 6e¯ ——> 2Cr3+ each Cr needs 3 electrons Step 4 Cr2O72- + 6e¯ + 14H+ ——> 2Cr3+ Step 5 Cr2O72- + 6e¯ + 14H+ ——> 2Cr3+ + 7H2O now balanced**BALANCING REDOX HALF EQUATIONS**Q. Balance the following half equations... Na —> Na+ Fe2+ —> Fe3+ I2 —> I¯ C2O42- —> CO2 H2O2 —> O2 H2O2 —> H2O NO3- —> NO NO3- —> NO2 SO42- —> SO2 REMINDER 1 Work out the formula of the species before and after the change; balance if required 2 Work out the oxidation state of the element before and after the change 3 Add electrons to one side of the equation so that the oxidation states balance 4 If the charges on all the species (ions and electrons) on either side of the equation do not balance then add sufficient H+ ions to one of the sides to balance the charges 5 If the equation still doesn’t balance, add sufficient water molecules to one side**BALANCING REDOX HALF EQUATIONS**Q. Balance the following half equations... Na —> Na++ e- Fe2+ —> Fe3+ + e- I2+ 2e- —> 2I¯ C2O42- —> 2CO2 + 2e- H2O2 —> O2 + 2H+ + 2e- H2O2 + 2H+ + 2e- —> 2H2O NO3-+ 4H+ + 3e- —> NO+ 2H2O NO3-+ 2H+ + e- —> NO2+ H2O SO42- + 4H+ + 2e- —> SO2 + 2H2O**COMBINING HALF EQUATIONS**A combination of two ionic half equations, one involving oxidation and the other reduction, produces a REDOX equation. The equations are balanced as follows... Step 1 Write out the two half equations Step 2 Multiply the equations so that the number of electrons in each is the same Step 3 Add the two equations and cancel out the electrons on either side Step 4 If necessary, cancel any other species which appear on both sides**COMBINING HALF EQUATIONS**A combination of two ionic half equations, one involving oxidation and the other reduction, produces a REDOX equation. The equations are balanced as follows... Step 1 Write out the two half equations Step 2 Multiply the equations so that the number of electrons in each is the same Step 3 Add the two equations and cancel out the electrons on either side Step 4 If necessary, cancel any other species which appear on both sides The reaction between manganate(VII) and iron(II)**COMBINING HALF EQUATIONS**A combination of two ionic half equations, one involving oxidation and the other reduction, produces a REDOX equation. The equations are balanced as follows... Step 1 Write out the two half equations Step 2 Multiply the equations so that the number of electrons in each is the same Step 3 Add the two equations and cancel out the electrons on either side Step 4 If necessary, cancel any other species which appear on both sides The reaction between manganate(VII) and iron(II) Step 1 Fe2+ ——> Fe3+ + e¯ Oxidation MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O Reduction**COMBINING HALF EQUATIONS**A combination of two ionic half equations, one involving oxidation and the other reduction, produces a REDOX equation. The equations are balanced as follows... Step 1 Write out the two half equations Step 2 Multiply the equations so that the number of electrons in each is the same Step 3 Add the two equations and cancel out the electrons on either side Step 4 If necessary, cancel any other species which appear on both sides The reaction between manganate(VII) and iron(II) Step 1 Fe2+ ——> Fe3+ + e¯ Oxidation MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O Reduction Step 2 5Fe2+ ——> 5Fe3+ + 5e¯ multiplied by 5 MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O multiplied by 1**COMBINING HALF EQUATIONS**A combination of two ionic half equations, one involving oxidation and the other reduction, produces a REDOX equation. The equations are balanced as follows... Step 1 Write out the two half equations Step 2 Multiply the equations so that the number of electrons in each is the same Step 3 Add the two equations and cancel out the electrons on either side Step 4 If necessary, cancel any other species which appear on both sides The reaction between manganate(VII) and iron(II) Step 1 Fe2+ ——> Fe3+ + e¯ Oxidation MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O Reduction Step 2 5Fe2+ ——> 5Fe3+ + 5e¯ multiplied by 5 MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O multiplied by 1 Step 3 MnO4¯ + 5e¯ + 8H+ + 5Fe2+ ——> Mn2+ + 4H2O + 5Fe3+ + 5e¯**COMBINING HALF EQUATIONS**A combination of two ionic half equations, one involving oxidation and the other reduction, produces a REDOX equation. The equations are balanced as follows... Step 1 Write out the two half equations Step 2 Multiply the equations so that the number of electrons in each is the same Step 3 Add the two equations and cancel out the electrons on either side Step 4 If necessary, cancel any other species which appear on both sides The reaction between manganate(VII) and iron(II) Step 1 Fe2+ ——> Fe3+ + e¯ Oxidation MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O Reduction Step 2 5Fe2+ ——> 5Fe3+ + 5e¯ multiplied by 5 MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O multiplied by 1 Step 3 MnO4¯ + 5e¯ + 8H+ + 5Fe2+ ——> Mn2+ + 4H2O + 5Fe3+ + 5e¯ Step 4 MnO4¯ + 8H+ + 5Fe2+ ——> Mn2+ + 4H2O + 5Fe3+**COMBINING HALF EQUATIONS**A combination of two ionic half equations, one involving oxidation and the other reduction, produces a REDOX equation. The equations are balanced as follows... Step 1 Write out the two half equations Step 2 Multiply the equations so that the number of electrons in each is the same Step 3 Add the two equations and cancel out the electrons on either side Step 4 If necessary, cancel any other species which appear on both sides The reaction between manganate(VII) and iron(II) Step 1 Fe2+ ——> Fe3+ + e¯ Oxidation MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O Reduction Step 2 5Fe2+ ——> 5Fe3+ + 5e¯ multiplied by 5 MnO4¯ + 5e¯ + 8H+ ——> Mn2+ + 4H2O multiplied by 1 Step 3 MnO4¯ + 5e¯ + 8H+ + 5Fe2+ ——> Mn2+ + 4H2O + 5Fe3+ + 5e¯ Step 4 MnO4¯ + 8H+ + 5Fe2+ ——> Mn2+ + 4H2O + 5Fe3+ SUMMARY**COMBINING HALF EQUATIONS**A combination of two ionic half equations, one involving oxidation and the other reduction, produces a REDOX equation. The equations are balanced as follows... Step 1 Write out the two half equations Step 2 Multiply the equations so that the number of electrons in each is the same Step 3 Add the two equations and cancel out the electrons on either side Step 4 If necessary, cancel any other species which appear on both sides Q.Construct balanced redox equations for the reactions between... Mg and H+ Cr2O72- and Fe2+ H2O2 and MnO4¯ C2O42- and MnO4¯ S2O32- and I2 Cr2O72- and I¯