Redox reaction – chemical reactions in which both oxidation and reduction occur simultaneously.

1) Oxidation

gain of oxygen, O2 by a substance loss of hydrogen, H2 from a substance a loss of electrons occurs when there is an increase in oxidation number

2) Reduction

loss of oxygen, O2 by a substance gain of hydrogen, H2 from a substance a gain of electrons occurs when there is an decrease in oxidation number

1. Oxidation and Reduction in Terms of Gain and Loss of Oxygen

2CuO (s) + C (s) –> 2Cu (s) + CO2 (g)

Reduction:CuO loses its oxygen to form copper. CuO (oxidising agent) is being reduced.

Oxidation:Carbon gains oxygen to form CO2. Carbon (reducing agent) is said to be oxidised.

PbO (s) + CO (g) –> Pb (s) + CO2 (g)

Reduction:PbO loses its oxygen to form lead. PbO (oxidising agent) is being reduced.

Oxidation:Carbon monoxide (CO) gains oxygen to form CO2. Carbon monoxide (reducing agent) is said to be oxidised.

H2 (g) + CuO (s) –> H2O (l) + Cu (s)

Reduction:CuO loses its oxygen to form copper. CuO (oxidising agent) is being reduced.

Oxidation:Hydrogen (H2) gains oxygen to form H2O. Hydrogen (reducing agent) is said to be oxidised.

2. Oxidation and Reduction in Terms of Gain and Loss of Hydrogen

H2S (g) + Cl2 (g) –> S (s) + 2HCl (g)

Reduction:Cl2 gains hydrogen to form hydrogen chloride. Cl2 (oxidising agent) is being reduced.

Oxidation:Hydrogen sulphide loses hydrogen to form sulphur. Hydrogen sulphide (reducing agent) is said to be oxidised.

2NH3 (g) + 3Br2 (g) –> N2 (g) + 6HBr (g)

Reduction:Bromine gains hydrogen to form hydrogen bromide. Br2 (oxidising agent) is being reduced.

Oxidation:Ammonia loses hydrogen to form nitrogen. Ammonia (reducing agent) is said to be oxidized.

Oxidation number– is the charge that the atom of the element would have if complete transfer of electron takes place.

IUPAC nomenclature – name inorganic compounds in order to avoid confusion due to elements have variable oxidation numbers.

(i) Free elements have an oxidation number of zero.

Na = 0Mg = 0C = 0H2 = 0Br2 = 0

(ii) Oxidation number of a simple monoatomic ions is the same as its charge.

Na+ ion = +1Mg2+ ion = +2O2- ion = -2Cl- ion = -1H+ ion = +1

(iii) Sum of the oxidation number for a neutral compound is zero.

CaH2

(+2) + 2(-1)= 0Sum of oxidation number is 0

Al2O3

2(+3) + 3(-2)= 0Sum of oxidation number is 0

Iodine, Bromine, Chlorine, Nitrogen, Oxygen, Fluorine

—> Electronegativity increase

Cl2O2(+1) + (-2)= 0Sum of oxidation number is 0.

(Chlorine, bromine and iodine usually have the oxidation number of -1 except when combine with a more electronegative element.)

HClO(+1) + (+1) + (-2)= 0Sum of oxidation number is 0.

(Chlorine, bromine and iodine usually have the oxidation number of -1 except when combine with a more electronegative element.)

(iv) Polyatomic ion, the sum of the oxidation numbers of all the atoms equals the charge on the ion.

SO4 2-

(+6) + 4 (-2)= +6 + (-8)= -2Sum of oxidation number is -2

Cr2O72-

2(+6) + 7(-2)= -2Sum of oxidation number is -2

(v) Calculating the oxidation numbers of elements in compounds or ions.

K2Cr2O7

2 (+1) + 2x + 7 (-2) = 0x = +6Oxidation number of chromium in K2Cr2O7 is +6

NO3-

x + 3(-2) = -1x = +5Oxidation number of nitrogen in NO3- is +5

Hydrogen peroxide, H2O2

2(+1) + 2x = 0x = -1Oxidation number of oxygen in H2O2 is -1 (and not -2)(Usually oxidation number for combined oxygen usually is -2 except in peroxides)

F2O2(-1) + x = 0x = +2Oxidation number of oxygen in F2O is +2 (and not -2)(Usually oxidation number for combined oxygen usually is -2 except in fluorine compounds)

NaH(+1) + x = 0x = -1Oxidation number of hydrogen in NaH is -1 (and not +1)(Usually oxidation number for combined hydrogen usually is +1 except in metal hydrides)

AlH3

(+3) + 3x = 0x = -1Oxidation number of hydrogen in AlH3 is -1 (and not +1)(Usually oxidation number for combined hydrogen usually is +1 except in metal hydrides)

MgH2

(+2) + 2x = 0x = -1Oxidation number of hydrogen in MgH2 is -1 (and not +1)(Usually oxidation number for combined hydrogen usually is +1 except in metal hydrides)

(vi) Some metals show different oxidation numbers.

Compound Oxidation number of manganeseMnSO4 +2MnO2 +4K2MnO4 +6KMnO4 +7Compound Oxidation number of chromiumK2CrO4 +6K2Cr2O7 +6

(vii) Usually non-metals have negative oxidation numbers but Cl, Br & I can have positive or negative oxidation number.

Compound Oxidation number of chlorineHCl -1

ClO2 +4HClO4 +7

Stay tune for the next installment in the “Oxidation and Reduction” series with focus on the difference between redox reaction and

non-redox reactions.

3. Oxidation and Reduction in Terms of Changes in Oxidation Numbers

Redox reactions – oxidation number of all elements change.

Rusting of iron, combustion, displacement of metal from its salt solution, displacement of halogen from its halide solution and electrolysis are redox reaction.

-10 …. -3  -2  -1  0  +1  +2  +3  …  +10

<———-  Reduction || Oxidation ———->

H2 (g) + CuO (s) –> H2O (l) + Cu (s)Hydrogen: 0 –> +1 (Oxidised to water & Hydrogen is a reducing agent)Copper oxide: +2 –> 0 (Reduced to copper & Copper oxide is a oxidising agent)

2Zn (s) + O2 (g) –> 2ZnO (s)Zinc: 0 –> +2 (Oxidised to zinc ion & Zinc is a reducing agent)Oxygen: 0 –> -2 (Reduced to oxide ion & Oxygen is an oxidising agent)

2Mg (s) + CO2 (g) –> 2MgO (s) + C (s)Magnesium: 0 –> +2 (Oxidised to magnesium ion & Magnesium is a reducing agent)Carbon dioxide: +4 –> 0 (Reduced to carbon & Carbon dioxide is an oxidising agent)

Br2 (l) + 2HI (aq) –> 2HBr (aq) + I2 (s)Hydroiodic acid / Hydrogen iodide: -1 –> 0 (Oxidised to iodine & Hydroiodic acid is a reducing agent)Bromine: 0 –> -1 (Reduced to hydrobromic acid & Bromine is a oxidising agent)

Non-redox reactions – oxidation number of all elements remain unchanged.

Precipitation, decomposition and neutralisation are not redox reaction (non-redox reaction)

Precipitation:

AgNO3 (aq) + NaCl (aq) –> AgCl (s) + NaNO3 (aq)+1 +5 3(-2)      +1  -1              +1  -1        +1 +5  3(-2)

No change in the oxidation numbers.

Decomposition:

ZnCO3 (s) –> ZnO (s) + CO2 (g)+2 +4  3(-2)    +2 -2       +4  2(-2)

No change in the oxidation numbers.

Neutralisation:

NaOH (aq) + HCl (aq) –> NaCl (aq) + H2O (l)+1 -2 +1          +1 -1             +1  -1             2(+1)  -2

No change in the oxidation numbers.

Revision time:

Questions to ask yourself at this point:

1. Do you understand what is a redox reaction?2. Can you write the two half reactions out of a redox reaction?3. Do you understand the concept of oxidation number?4. Give three examples of an oxidising agent and the example of the reaction involved.5. Give three examples of a reducing agent and the example of the reaction involved.6. Can you differentiate a redox reaction with a non-redox reaction?7. What is the characteristics of a non-redox reaction?

4. Oxidation and Reduction in Terms of Electron Transfer

2I- (aq) –> I2 (aq) + 2eOxidation: Iodide ion, I- is a reducing agent because it donates/loses electrons to become I2.

Br2 + 2e –> 2Br- (aq)Reduction: Bromine water, Br2 is an oxidising agent because it receives/accepts electrons from I- to form bromide ions, Br-.

–> Overall reaction: 2I- (aq) + Br2 –> I2 (aq) + 2Br- (aq)

Conversion of Fe 2+ Ions to Fe 3+ Ions and Vice Versa

A) Common oxidising agent (change Fe2+ ions to Fe3+ ions):

bromine, Br2

chlorine, Cl2

hydrogen peroxide, H2O2

concentrated nitric acid, HNO3

acidified potassium manganate(VII), KMnO4 solution acidified potassium dichromate(VI), K2Cr2O7 solution

Fe2+ (aq) –> Fe3+ (aq) + eOxidation: Iron(II) ion, Fe2+ is a reducing agent because it donates/loses one electron to become Fe3+.

Br2 (aq) + 2e –> 2Br- (aq)Reduction: Bromine water, Br2 is an oxidising agent because it receives/accepts electrons from Fe2+ to form bromide ions, Br-.

–> Observation: iron(II) sulphate solution changes colour from pale green to yellowish-brown.–> Overall reaction: 2Fe2+ (aq) + Br2 (aq) –> 2Fe3+ (aq) +2Br- (aq)

B) Common reducing agent (change Fe3+ ions to Fe2+ions):

zinc powder, Zn aluminium, Al Magnesium, Mg Calcium, Ca Sulphur dioxide, SO2

Hydrogen sulphide, H2S Sodium sulphide solution, Na2SO3

Tin(II) chloride solution, SnCl2

Potassium iodide, KI

Zn (s) –> Zn2+ (aq) + 2eOxidation: Zinc powder, Zn is a reducing agent because it donates/loses two electrons to form zinc ions, Zn2+.

Fe3+ (aq) + e –> Fe2+ (aq)Reduction: Iron(III) ion, Fe3+ is an oxidising agent because it receives/accepts one electron to become Fe2+.

–> Observation: iron(III) sulphate solution changes colour from yellowish-brown to pale green.–> Overall reaction: 2Fe3+ (aq) + Zn (aq) –> 2Fe2+ (aq) + Zn2+ (aq)

C) Investigate the presence of iron(II) and iron(III) ions

Reagent Ions ObservationsNaOH solution / NH3 solution Fe2+ Green precipitate,insoluble in excess

alkaliNaOH solution / NH3 solution Fe3+ Brown precipitate,insoluble in excess

alkaliPotassium hexacyanoferrate(II) solution

Fe2+ Light blue precipitate

Potassium hexacyanoferrate(II) solution

Fe3+ Dark blue precipitate

Potassium hexacyanoferrate(III) solution

Fe2+ Dark blue precipitate

Potassium hexacyanoferrate(III) solution

Fe3+ Greenish-brown solution

Potasium / Ammonium thiocyanate solution

Fe2+ Pale red colouration

Potasium / Ammonium thiocyanate solution

Fe3+ Blood-red colouration

Displacement of Metals from Their Salt Solution

More electropositive metal is more reactive metal and it will displace a less reactive metal from the solution.

Reactivity Series

Metal EquationPotassium K (s) –> K+ (aq) + eSodium Na (s) –> Na+ (aq) + eCalcium Ca (s) –> Ca2+ (aq) + 2eMagnesium Mg (s) –> Mg2+ (aq) + 2eAluminium Al (s) –> Al3+ (aq) + 3eZinc Zn (s) –> Zn2+ (aq) + 2eIron Fe (s) –> Fe2+ (aq) + 2eTin Sn (s) –> Sn2+ (aq) + 2eLead Pb (s) –> Pb2+ (aq) + 2eHydrogen H (s) –> H+ (aq) + eCopper Cu (s) –> Cu2+ (aq) + 2eSilver Ag (s) –> Ag+ (aq) + e

Going up the table:

1. Tendency of metals to form positive ions increases.2. Electropositivity increases .3. Strength as reducing agent increases.

(The focus should be on the strength of the reducing agent)

Going down the table:

1. Tendency of metals to receive electrons increases.2. Electropositivity decreases .3. Strength as oxidising agent increases.

(The focus should be on the strength of the oxidising agent)

a) Magnesium ribbon + Lead(II) nitrate solution

Mg (s) –> Mg2+ (aq) + 2eOxidation: Magnesium ribbon, Mg is a reducing agent because it is more electropositive than lead.

Pb2+ (aq) + 2e –> Pb (s)Reduction: Lead(II) ion, Pb2+ is an oxidising agent because it receives/accepts two electrons to become lead.

Observation:i) Deposition of dark brown solid (lead) on the magnesium ribbon.ii) The magnesium ribbon dissolves.

Overall reaction: Pb2+ (aq) + Mg (s) –> Pb (s) + Mg2+ (aq)

b) Zinc strip + Copper(II) sulphate solution

Zn (s) –> Zn2+ (aq) + 2eOxidation: Zinc strip, Zn is a reducing agent because it is more electropositive than copper.

Cu2+ (aq) + 2e –> Cu (s)Reduction: Copper(II) ion, Cu2+ is an oxidising agent because it receives/accepts two electrons to become copper.

Observation:i) Deposition of brown solid (copper) on the zinc strip.ii) The blue colour of CuSO4 becomes paler / The blue solution turns to colourless.iii) The zinc strip dissolves.

Overall reaction: Cu2+ (aq) + Zn (s) –> Cu (s) + Zn2+ (aq)

c) Copper strip + Magnesium sulphate solution

No changes.Magnesium is more electropositive than copper.

Observation:No changes.There is no redox reaction occurs.

Mnemonic method for memorising the reactivity series

Students typically will form sentences (sometimes naughty ones) to memorise the series in a mnemonic way. Berry Berry Easy would like to share with you some famous ones such as:

Paddy Still Could Marry A Zulu In The Lovely Honolulu Causing Many Strange Gazes.Potassium Sodium Calcium Magnesium Aluminium Zinc Iron Tin Lead Hydrogen Copper Mercury Silver Gold(In the above statement, there is addition of Mercury and Gold. No harm memorising those too)

Displacement of Halogen (Group 17) from Its Halide Solution Halogen – elements in Group 17 of the Periodic Table Example: chlorine, bromine and iodine. Halogen Identification

Halogen Colour in (conc.) aq. solution

Colour in (dilute) aq. solution

Colour in 1,1,1-trichloroethane

Iodine Reddish-brown Yellow PurpleBromine Brown Yellow BrownChlorine Pale yellow Colourless Colourless

Strength of oxidising agent in halogen Cl2, Br2. I2

<——— Oxidising power increases

Halide / Halogen Chlorine Bromine IodinePotassium chloride

- No changes No changes

Potassium bromide

Chlorine displace bromine from KBr solution

- No changes

Potassium iodide Chlorine displace iodine from KI solution

Bromine displace iodine from KI solution

-

. Redox Reactions by the Transfer of Electrons at a Distance

Set I

Reducing agent Oxidising agent Test on the solution in the reducing agent arm of U-tube

Iron(II) sulphate, FeSO4 solution

Acidified potassium dichromate(VI), K2Cr2O7

solution

Add a few drops of potassium thiocyanate, KSCN solution

Observation InferenceThe electrode in the iron(II) sulphate, FeSO4 solution acts as the negative terminal, whereas the electrode in the acidified potassium dichromate(VI), K2Cr2O7 solution acts as the positive terminal.

Electrons flow from iron(II) sulphate, FeSO4 solution to acidified potassium dichromate(VI), K2Cr2O7 solution

Iron(II) sulphate solution changes from pale green to yellow/brown. It gives blood-red colouration with potassium thiocyanate solution (KSCN)

Iron(III) ions are present. Iron(II) ions are oxidised to  iron(III) ions.

Acidified potassium dichromate(VI), K2Cr2O7

solution changes colour from orange to green.Dichromate(VI) ions are reduced to chromium(III) ions.

Oxidation half-equation: Fe2+(aq) –> Fe3+(aq) + e Reduction half-equation: Cr2O7

2-(aq) + 14H+(aq) + 6e –> 2Cr3+(aq) + 7H2O(l) Overall reaction: Cr2O7

2-(aq) + 6Fe2+(aq) 14H+(aq) –> 2Cr3+(aq) + 6Fe3+(aq) + 7H2O(l)

Set II

Reducing agent Oxidising agent Test on the solution in the reducing agent arm of U-tube

Iron(II) sulphate, FeSO4 solution

Acidified manganate(VII), KMnO4 solution

Add sodium hydroxide, NaOH solution

Observation InferenceThe electrode in the iron(II) sulphate, FeSO4 Electrons flow from iron(II)

solution acts as the negative terminal, whereas the electrode in the acidified potassium manganate(VII), KMnO4 solution acts as the positive terminal.

sulphate, FeSO4 solution to acidified potassium manganate(VII), KMnO4 solution

Iron(II) sulphate solution changes from pale green to yellow/brown. It formed a brown precipitate when the brown solution is tested with sodium hydroxide solution (NaOH)

Iron(III) ions are present. Iron(II) ions are oxidised to  iron(III) ions.

Purple acidified manganate(VII), KMnO4 solution turns colourless.

Manganate(VII) ions are reduced to manganese(II) ions.

Oxidation half-equation: Fe2+(aq) –> Fe3+(aq) + e Reduction half-equation: MnO4

-(aq) + 8H+(aq) + 5e –> Mn2+(aq) + 4H2O(l) Overall reaction: MnO4

-(aq) + 5Fe2+(aq) + 8H+(aq) –> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)

Set III

Reducing agent Oxidising agent Test on the solution in the reducing agent arm of U-tube

Potassium iodide, KI solution

Bromine water, Br2 Add a few drops of starch solution

Observation InferenceThe electrode in the potassium iodide, KI solution acts as the negative terminal, whereas the electrode in the bromine water acts as the positive terminal.

Electrons flow from potassium iodide, KI solution to bromine water, Br2 (aq).

Colourless potassium iodide solution turns brown. It formed a dark blue colouration when the brown solution is tested with starch solution.

Iodine is present.Iodide ions have oxidised to iodine.

Brown bromine water turns colourless. Bromines are reduced to bromide ion.

Oxidation half-equation: 2I-(aq) –> I2(aq) + 2e Reduction half-equation: Br2(aq) + 2e –> 2Br-(aq) Overall reaction: Br2(aq) + 2I-(aq) –> 2Br-(aq) + I2(aq)

Other pairs of reducing agent and oxidising agent

Reducing agent Oxidising agentPotassium iodide,KI solution Iron(III) sulphate,Fe2(SO4)3 solutionPotassium iodide,KI solution Acidified potassium dichromate(VI),K2Cr2O7 solutionPotassium bromide,KBr solution Chlorine, Cl2 water

Rusting as a Redox Reaction

Rust / hydrated iron(III) oxide, Fe2O3•xH2O – formed slowly at the surface of iron when it exposed to the damp air.

Rusting – a redox reaction that take places between iron and oxygen to form hydrated iron(III) oxide and this is a slow reaction.4Fe(s) + 3O2(g) + 2xH2O(l) –> Fe2O3•xH2O(s)

Corrosion – a redox reaction that take places between a metal and the gases in air. Metal is oxidised to form an oxide layer on the surface. Metal atoms lose electrons to form positive ions.

1. Group 1 metals are very reactive.2. Metals are exposed to air will corrode rapidly and become tarnished.3. Aluminium , lead and zinc corrode rapidly in the air and forms an oxide layer. The oxide layer is hard, non-porous,

impermeable and difficult to crack. This protects the aluminium, lead and zinc below it from further corrosion.

Example: Corrosion of metal.Zn(s) –> Zn2+(aq) + 2eCu(s) –> Cu2+(aq) + 2e

K, Na, Ca, Mg, Al, Zn, Fe, Sn, Pb, Cu, Hg, Ag, Au<————Tendency of metal to corrode increases.

Process of Rusting of Iron

1. Anode (negative terminal) – concentration of oxygen is lower and iron rust (oxidation process) to form iron(II) ions:Fe(s) –> Fe2+(aq) + 2e

2. Cathode (positive terminal) – concentration of oxygen is higher and oxygen gains electrons that reduced to hydroxide ions:O2(g) + 2H2O(l) + 4e –> 4OH-(aq)

3. Fe2+ ions and OH- ions combine to form iron(II) hydroxide, Fe(OH)2

4. Oxygen further oxidises iron(II) hydroxide, Fe(OH)2 to hydrated iron(III) oxide, Fe2O3•xH2O.2Fe(OH)2(s) –> Fe2O3•xH2O(s)

The Effect of Other Metals on the Rusting of Iron

Potassium hexacyanoferrate(III), K3Fe(NO)6 is used to detect Fe2+ ions(produces dark blue colour in the presence of Fe2+).

Phenolphthalein is used to detect OH- ions(produces pink colour in the presence of OH-).

Test tube Observation ReactionFe only Intensity of blue colour

is low.Oxidation:Fe(s) –> Fe2+(aq) + 2e

Control Pink colour is not present.

Reduction:O2(g) + 2H2O(l) + e –> 4OH-(aq)OH- ions react with Fe2+ ions to form Fe(OH)2.

Inference Fe nail rusts a little.

Test tube Observation ReactionFe & Mg Blue colour is not

present.Oxidation:Mg(s) –> Mg2+(aq) + 2e

Intensity of pink colour is very high.

Reduction:O2(g) + 2H2O(l) + e –> 4OH-(aq)

Inference Mg is corroded and Fe nail does not rust.

i) Fe act as the (+) terminal (cathode)ii) Mg act as the (-) terminal (anode)

Test tube Observation ReactionFe & Zn Blue colour is not

present.Oxidation:Zn(s) –> Zn2+(aq) + 2e

Intensity of pink colour is high.

Reduction:O2(g) + 2H2O(l) + e –> 4OH-(aq)

Inference Zn is corroded and Fe nail does not rust.

i) Fe act as the (+) terminal (cathode)ii) Zn act as the (-) terminal (anode)

Test tube Observation ReactionFe & Sn Intensity of blue colour

is high.Oxidation:Fe (s) –> Fe2+(aq) + 2e

Pink colour is not present.

Reduction:O2(g) + 2H2O(l) + e –> 4OH-(aq)OH- ions react with Fe2+ ions to form Fe(OH)2.

Inference Fe nail rusts quickly (high rate).

i) Sn act as the (+) terminal (cathode)ii) Fe act as the (-) terminal (anode)

Test tube Observation ReactionFe & Cu Intensity of blue colour

is very high.Oxidation:Fe (s) –> Fe2+(aq) + 2e

Pink colour is not present.

Reduction:O2(g) + 2H2O(l) + e –> 4OH-(aq)OH- ions react with Fe2+ ions to form Fe(OH)2.

Inference Fe nail rusts very quickly (the highest rate).

i) Cu act as the (+) terminal (cathode)ii) Fe act as the (-) terminal (anode)

Iron nail does not rust if it has contacted with more electropositive metals (Mg and Zn). Iron nail rusts quickly if it has contacted with less electropositive metals (Sn and Cu).

Prevention of Rusting of Iron

The rate of rusting of iron decreases if the iron (Fe) in contact with any of these metals: K, Na, Ca, Mg, Al and Zn.The rate of rusting of iron increases if the iron (Fe) in contact

with any of these metals: Sn, Pb, Cu, Hg, Ag and Au. a strong electrolyte (salt and acid) is present.

Ways Used for Prevention of Rusting

1. Painting – Protect iron surface (prevent from contacting with air and water)2. Coat with plastic – Used in metal netting3. Apply oil and grease – Protective coating for machine part4. Alloying the iron – Alloying the iron with 18% chromium and 8% nickel that provide a protective oxide coating.5. a) Tin plating (less electropositive metal) – Cans of food (iron) is covered with a thin layer of tin to provide a protective

oxide coating to the cans.b) Chrome plating

6. Cathodic protection / Electrical protection (more electropositive metal)a) Galvanising (coat with zinc metal) – Zinc layer provides a protective oxide coating and zinc is oxidized instead of iron. Iron cannot form ions, so it will not rust.b) Sacrificial protection – Blocks of magnesium are attached at the intervals of the water piping system & zinc bars are attached to the part of the ship submerged in sea water.

The Reactivity Series of Metals and Its Application

1. Metal form metal oxides when burnt in air (excess).Metal + Oxygen –> Metal oxideExample: 2Zn(s) + O2(g) –> 2ZnO(s)

2. The more reactive a metal is, the more vigorously it burns in oxygen.

Reactivity of MetalsK, Na, Ca, Mg, Al, Zn, Fe, Sn, Pb, Cu, Hg, Ag, Au<——– increase in reactivity of metals

3. Reactivity of Metals with Oxygen.

Metal Observation Inference EquationMagnesium (Mg) Burns vigorously

with a very brilliant white flame.The residue is white when hot and cold.

The reactivity of Mg towards O2 is very high.Magnesium oxide is formed.

2Mg(s) + O2(g)–> 2MgO(s)

Zinc (Zn) Burns quickly with a bright flame.The residue is yellow when hot and white when cold.

The reactivity of Zn towards O2 is high.Zinc oxide is formed.

2Zn(s) + O2(g)–> 2ZnO(s)

Iron (Fe) Glows very brightly.The residue is reddish-brown when hot and cold.

The reactivity of Fe towards O2 is medium.Iron(III) oxide is formed.

2Fe(s) + O2(g)–> 2Fe2O3(s)

Lead (Pb) Glows brightly.The residue is brown when hot and yellow when cold.

The reactivity of Pb towards O2 is low.Lead(II) oxide is formed.

2Pb(s) + O2(g)–> 2PbO(s)

Copper (Cu) Glows faintly.The residue is black when hot and cold.

The reactivity of Cu towards O2 is very low.Copper(II) oxide is formed.

2Cu(s) + O2(g)–> 2CuO(s)

Glass wool – prevents metal powder mixes with potassium manganate(VII) Solid potassium manganate(VII) – liberates oxygen gas when it is heated / decomposed.

2KMnO4(s) —-> K2MnO4(s) + MnO2(s) + O2(g)heat

Other than potassium manganate(VII),

- solid potassium chlorate(V) with manganese(IV) oxide as a catalyst.

MnO2

2KClO3(s) —-> KCl(s) + 3O2(g)heat

- solid potassium nitrate

2KNO3(s) —-> KNO2(s) + O2(g)heat

4. Position of Carbon in the Reactivity Series of Metals

Reactivity SeriesK, Na, Ca, Mg, Al, C, Zn, Fe, Sn, Pb, Cu, Hg, Ag, Au<——– increase in reactivity

a) Metal + Carbon dioxide –> Metal oxide + Carbon

Any metal above carbon in the reactivity series can displace oxygen from carbon dioxide.

Example: 2Mg(s) + CO2(g) –> 2MgO(s) + C(s)

Mg: Reducing agentCO2: Oxidising agentMgO: White residueC: Black spots

–> Therefore, magnesium is more reactive than carbon.(If the metal is less reactive than carbon, the metal is unable to remove oxygen from carbon dioxide.)

b) Carbon + Metal oxide –> Carbon dioxide + Metal

Any metal below carbon in the reactivity series can displace carbon from its oxide.

Example: C(s) + 2ZnO(s) –> 2Zn(s) + CO2(g)

C: Reducing agentZnO: Oxidising agentZn: Grey residue

–> Therefore, zinc is less reactive than carbon.(If carbon is less reactive than the metal, the carbon is unable to remove oxygen from metal oxide.)

Position of Hydrogen in the Reactivity Series of Metals

Reactivity SeriesK, Na, Ca, Mg, Al, C, Zn, H, Fe, Sn, Pb, Cu, Hg, Ag, Au<——– increase in reactivity

Metal oxide + Hydrogen –> Metal + Water

Any metal below hydrogen in the reactivity series, hydrogen will reduce the oxide of metal to metal.

Example 1:

CuO(s) + H2(g) –> Cu(s) + H2O(l) Observation : Burns quickly with a bright flame. The black solid turns brown solid. H2: Reducing agent CuO: Oxidising agent Hydrogen is more reactive than copper.

Example 2:

PbO(s) + H2(g) –> Pb(s) + H2O(l) Observation : Burns with a bright flame. The yellow solid turns grey solid. H2: Reducing agent PbO: Oxidising agent Hydrogen is more reactive than lead.

Example 3:

Fe2O3(s) + 3H2(g) –> 2Fe(s) + 3H2O(l)

Observation : Glows dimly. The reddish-brown solid turns grey solid. H2: Reducing agent Fe2O3: Oxidising agent Hydrogen is more reactive than iron.

Example 4:

ZnO(s) + H2(g) –> no reaction Observation : No glow is observed. It turns yellow when hot and white when cold. Hydrogen is unable to reduce zinc oxide. Hydrogen is less reactive than zinc.

Redox Reactions in Electrolytic Cell and Chemical Cell

Similarities

redox reaction. Anode: oxidation Cathode: reduction Electrons flow from anode to cathode in the external circuit

Differences

Differences Electrolytic Cell (Electrolysis)

Chemical Cell / Voltaic Cell

Structure With electrical supply. No electrical supply.Electrodes Can be the same or

difference metal (graphite or platinum).

Must be two different metals.

Flows of electrons From anode to cathode through external circuit.

From more electropositive metal to less electropositive metal through external circuit.

Transformation of energy

Electrical energy to chemical energy.

Chemical energy to electrical energy.

At positive terminal Anode.Oxidation occurs. Anions release electrons at the anode.

Cathode.Reduction occurs.

Oxidising agent gain electrons.

At negative terminal Cathode.Reduction occurs.

Cations gain electrons from the cathode.

Anode.Oxidation occurs.

Reducing agent releases electrons.

1) Redox Reactions in Electrolytic Cell

Example 1: Electrolysis of molten zinc chloride

Electrodes: Carbon Ions present: Cl- and Zn2+

Anode : Oxidation / 2Cl-(l) –> Cl2(g) + 2e / Cl- ions act as reducing agent. Cathode : Reduction / Zn2+(l) + 2e –> Zn(s) / Zn2+ ions act as oxidising agent.

Example 2: Electrolysis of copper(II) sulphate solution.

Electrodes: Carbon Ions present: Cu2+, SO4

2-, H+, OH-

OH- ions are discharged because OH- ion is below SO42- ion in electrochemistry series.

Anode: Oxidation / 4OH-(aq) –> O2(g) + 2H2O(l) + 4e / Oxygen gas is liberated. Cu2+ ions are discharged because Cu2+ ion is below H+ ion in electrochemistry series.

Cathode: Reduction / Cu2+(aq) + 2e –> Cu(s) / Cu2+ ions are reduced to copper metal (brown layer formed). –> Overall equation: Cu2+(aq) + 4OH-(aq) –> O2(g) + 2H2O(l) + Cu(s)

Example 3: Electrolysis of copper(II) sulphate solution.

Electrodes: Copper Ions present: Cu2+, SO4

2-, H+, OH-

OH- ions  and SO42- ion are not discharged.

Anode: Oxidation / Cu(s) –> Cu2+(aq) + 2e / Copper anode (electrode) dissolves. Cu2+ ions are discharged because Cu2+ ion is below H+ ion in electrochemistry series.

Cathode: Reduction / Cu2+(aq) + 2e –> Cu(s) / Cu2+ ions are reduced to copper metal.

Example 4: Electrolysis of concentrated sodium chloride solution.

Electrodes: Carbon Ions present: Na+, Cl-, H+, OH-

Cl- ions are discharged because of the higher concentration. (Concentration of Cl- ion is high, the ion is selectively discharged rather than the OH- ion, the one that is placed below the electrochemical series.)Anode: Oxidation / 2Cl-(aq) –> Cl2(g) + 2e / Chlorine gas (green gas with choking smell) is liberated.

H+ ions are discharged because Na+ ion is below H+ ion in electrochemistry series. (H+ ions and Na+ ion are placed very far apart in the electrochemical series, the concentration factor becomes unimportant.)Cathode: Reduction / 2H+(aq) + 2e –> H2(g) / Hydrogen gas is liberated.

–> Overall equation: 2Cl-(aq) + 2H+(aq) –> Cl2(g) + H2(g)

2) Redox Reactions in Chemical Cell

Example 1: Daniel cell

Anode (negative terminal): Oxidation / Zinc strip immerses in zinc sulphate solution.Zn(s) –> Zn2+(aq) + 2e / Zinc strip becomes thinner.

Cathode (positive terminal): Reduction / Copper strip immerses in copper(II) sulphate solution.Cu2+(aq) + 2e –> Cu(s) / A brown layer formed around copper strip. / Concentration Cu2+ ions decreases cause the intensity blue colour of solution decreases.

Zinc is more electropositive than copper. Electrons are flowed from zinc strip to copper strip through the external circuit. (Note: Conventionally, electrons flow in the opposite direction of electrical current).

–> Overall equation:  Zn(s) + Cu2+(aq) –> Zn2+(aq) + Cu(s)

More Chemical Cells

1. Primary cells – are not rechargeable and can be used only once.

2. Secondary cells – are rechargeable when cells are exhausted and can be reused again.

A. Dry Cell

Anode (-): Zinc / Zn(s) –> Zn2+(aq) + 2e / reducing agent Cathode (+): Graphite (carbon) rod / 2NH4

+(aq) + 2e –> 2NH3(g) + H2(g) / oxidising agent Electrolyte: Moist paste of ammonium chloride, zinc chloride and a little water. Overall reaction: Zn(s) + 2NH4

+(aq) –> Zn2+(aq) + 2NH3(g) + H2(g) Uses: touchlight, toys, clock, remote control and radio.

B. Alkaline Cell

Anode (-): Zinc / Zn(s) –> Zn2+(aq) + 2e / reducing agent

Cathode (+): Manganese(IV) oxide / 2MnO2(s) + H2O(l) +2e –> Mn2O3(s) + 2OH-(aq) / oxidising agent Electrolyte: Potassium hydroxide paste. Overall reaction: Zn(s) + 2MnO2(s) + H2O(l) –> Zn2+(aq) + Mn2O3(s) + 2OH-(aq) Heavy use and longer shelf life. Zinc corrodes more slowly. Higher power. More stable current and voltage.

C. Mercury Cell

Anode (-): Zinc / Zn(s) –> Zn2+(aq) + 2e / reducing agent Cathode (+): Mercury(II) oxide / Hg2+(aq) + 2e –> Hg(l) / oxidising agent Electrolyte: Potassium hydroxide paste. Overall reaction: Zn(s) + Hg2+(aq) –> Zn2+(aq) + Hg(l) Danger to the environment and mercury need to recycle. Uses: Watches, camera and small devices.

D. Lead-acid Accumulator

Anode (-): Lead / Pb(s) + SO42-(aq) –> PbSO4(s) + 2e / reducing agent

Cathode (+): Lead(IV) oxide / PbO2(s) + 4H+(aq) + SO42-(aq) + 2e –> PbSO4(s) + 2H2O(l) / oxidising agent

Electrolyte: Sulphuric acid. Overall reaction: / Pb(s) + PbO2(s) + 4H+(aq) + 2 SO4

2-(aq) –> 2PbSO4(s) + 2H2O(l) Uses: Automobiles.

E. Nickel-Cadmium Cell

Anode (-): Cadmium / Cd(s) + 2OH-(aq) –> Cd(OH)2(s) + 2e / reducing agent Cathode (+): Nickel(IV) oxide / NiO2(s) + 2H2O(l) + 2e –> Ni(OH)2(s) + 2OH-(aq) / oxidising agent Electrolyte: Porous separator soaked in potassium hydroxide solution. Overall reaction: Cd(s) + NiO2(s) + 2H2O(l) –> Cd(OH)2(s) + Ni(OH)2(s) Suffer from memory effect – hold less charge. Toxic heavy metal. Expensive. Uses: Toys, laptops, and mobile phones.

F. Rechargeable Chemical Cell

i) Nickel-metal hydride (NiMH)

Anode (-): hydrogen-absorbing alloy. Cathode (+): Nickel(IV) oxide. Contains rare earth elements such as titanium, vanadium, zirconium, cobalt, manganese and aluminium that are more

environmentally friendly. Higher capacity than NiCd. Higher self-discharge rate. Uses: digital cameras and mobile phones.

ii) Lithium-ion (Li-Ion)

Smaller and lighter. Anode (-): Carbon. Cathode (+): Metal oxide (cobalt oxide / manganese oxide). Electrolyte: Lithium salt in an organic solvent (ether). Inflammable and can easily explode when exposed to high temperature. Uses: Portable electronic.

iii) Lithium-polymer (Li-Poly)

Very small, thin and light. Anode (-): Carbon. Cathode (+): Metal oxide. Electrolyte: Lithium salt in a solid polymer composite (polyacrylonitrile). Not flammable. Uses: MP3, PDAs and laptops.

G. Other Chemical Cells

i) Fuel Cells

Anode (-): Fuel (hydrogen / hydrocarbon / alcohol). Cathode (+): Oxygen. Non-polluting product. Uses: space vehicles and military applications.

ii) Solar Cells

Made of semiconductor materials (crystalline silicon). Solar energy converted to electric energy. Non-polluting product. High cost. Uses: space satellites, irrigation pumps, calculator and telecommunications.