More Chemical Cells

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) + SO₄^2-(aq) –> PbSO₄(s) + 2e / reducing agent
• Cathode (+): Lead(IV) oxide / PbO₂(s) + 4H⁺(aq) + SO₄^2-(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.

Read More

Position of Hydrogen in the Reactivity Series of Metals

Position of Hydrogen in the Reactivity Series of Metals
Reactivity Series
K, 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 SO₄^2- 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 SO₄^2- 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)

Read More

The Reactivity Series of Metals and Its Application

The Reactivity Series of Metals and Its Application
1. Metal form metal oxides when burnt in air (excess).
Metal + Oxygen –> Metal oxide
Example: 2Zn(s) + O₂(g) –> 2ZnO(s)
2. The more reactive a metal is, the more vigorously it burns in oxygen.
Reactivity of Metals
K, 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	Equation
Magnesium (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 Series
K, 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 agent
CO₂: Oxidising agent
MgO: White residue
C: 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 agent
ZnO: Oxidising agent
Zn: 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.)

Read More

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) + 2e
Cu(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	Reaction
Fe 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	Reaction
Fe & 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	Reaction
Fe & 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	Reaction
Fe & 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	Reaction
Fe & 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 netting
3. Apply oil and grease – Protective coating for machine part
4. 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.

Read More

Redox Reactions by the Transfer of Electrons at a Distance

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		Inference
The 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		Inference
The electrode in the iron(II) sulphate, FeSO4 solution acts as the negative terminal, whereas the electrode in the acidified potassium manganate(VII), KMnO4 solution acts as the positive terminal.		Electrons flow from iron(II) 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		Inference
The 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 agent
Potassium iodide,KI solution	Iron(III) sulphate,Fe2(SO4)3 solution
Potassium iodide,KI solution	Acidified potassium dichromate(VI),K2Cr2O7 solution
Potassium bromide,KBr solution	Chlorine, Cl2 water

Read More

The reactivity series

Rusting of Iron Nails

Despite the confusion in some revision books, the electrochemical and reactivity series is not the same.
The reactivity series is about:

• List of metals (sometimes with hydrogen and carbon added as baseline) arranged by the ability to reduce other chemicals in non-specific way.
• The reactivity is dependent on the conditions of the reactions and all are relative.
• So the order will be slightly different between various books and research. (However, generally it should be the same)
• Reactivity series
  K, *Na, *Ca,  Mg, Al, Zn, Fe, Sn, Pb, Cu, Hg, Ag, Au
  <—- increase in reactivity of metals.
  Do take notes with the *, the position of the calcium and sodium are different. In electrochemical series, calcium is more electropositive than sodium but not for the reactivity series.

This is in contract with the electrochemical series :

• List of reducing agents arranged by the way of negative potential in an electrochemical cell
• The conditions are pressure of 1 atm, temperature of 298 K and molar concentration of 1.o M
• The electrochemical series / electropositivity of metal
  

  Equilibrium	Eɵ (volts)
  Li+ (aq) + e <—> Li (s)	- 3.03
  K+ (aq) + e <—> K (s)	- 2.92
  *Ca2+ (aq) + 2e <—> Ca (s)	- 2.87
  *Na+ (aq) + e <—> Na (s)	- 2.71
  Mg2+ (aq) + 2e <—> Mg (s)	- 2.37
  Al3+ (aq) + 3e <—> Al (s)	- 1.66
  Zn2+ (aq) + 2e <—> Zn (s)	- 0.76
  Fe2+ (aq) + 2e <—> Fe (s)	- 0.44
  Pb2+ (aq) + 2e <—> Pb (s)	- 0.13
  2H+ (aq) + 2e <—> H2 (g)	0
  Cu2+ (aq) + 2e <—> Cu (s)	+ 0.34
  Ag+ (aq) + e <—> Ag (s)	+ 0.80
  Au3+ (aq) + 3e <—> Au (s)	+ 1.50

So, more IMPORTANTLY, the SPM syllabus focuses on the reactivity instead of electrochemical potentials, so do use the reactivity series instead of the electrochemical series.
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	Purple
Bromine	Brown	Yellow	Brown
Chlorine	Pale yellow	Colourless	Colourless

Strength of oxidising agent in halogen
Cl₂, Br₂. I₂
<——— Oxidising power increases

Halide / Halogen	Chlorine	Bromine	Iodine
Potassium 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	-

Read More

Electron Transfer

Rusting of Iron Nails

Oxidation and Reduction in Terms of Electron Transfer
2I^- (aq) –> I₂ (aq) + 2e
Oxidation: 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 Fe²⁺ Ions to Fe³⁺ 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) + e
Oxidation: 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

Zn (s) –> Zn²⁺ (aq) + 2e
Oxidation: 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	Observations
NaOH solution / NH3 solution	Fe2+	Green precipitate,insoluble in excess alkali
NaOH solution / NH3 solution	Fe3+	Brown precipitate,insoluble in excess alkali
Potassium 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

Read More

Changes in Oxidation Numbers Redox reactions

Rusting of Iron Nails

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.

Read More

Oxidation Number

Rusting of Iron Nails

Oxidation Number – is the charge that the atom of the element would have if complete transfer of electron takes place.
Oxidation number
(i) Free elements have an oxidation number of zero.
Na = 0
Mg = 0
C = 0
H₂ = 0
Br₂ = 0
(ii) Oxidation number of a simple monoatomic ions is the same as its charge.
Na⁺ ion = +1
Mg²⁺ ion = +2
O^2- ion = -2
Cl^- ion = -1
H⁺ ion = +1
(iii) Sum of the oxidation number for a neutral compound is zero.
CaH₂
(+2) + 2(-1)
= 0
Sum of oxidation number is 0
Al₂O₃
2(+3) + 3(-2)
= 0
Sum of oxidation number is 0
Iodine, Bromine, Chlorine, Nitrogen, Oxygen, Fluorine
—> Electronegativity increase
Cl₂O
2(+1) + (-2)
= 0
Sum 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)
= 0
Sum 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)
= -2
Sum of oxidation number is -2
Cr₂O₇^2-
2(+6) + 7(-2)
= -2
Sum of oxidation number is -2
(v) Calculating the oxidation numbers of elements in compounds or ions.
K₂Cr₂O₇
2 (+1) + 2x + 7 (-2) = 0
x = +6
Oxidation number of chromium in K₂Cr₂O₇ is +6
NO^3-
x + 3(-2) = -1
x = +5
Oxidation number of nitrogen in NO^3- is +5
Hydrogen peroxide, H₂O₂
2(+1) + 2x = 0
x = -1
Oxidation number of oxygen in H₂O₂ is -1 (and not -2)
(Usually oxidation number for combined oxygen usually is -2 except in peroxides)
F₂O
2(-1) + x = 0
x = +2
Oxidation 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 = 0
x = -1
Oxidation 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 = 0
x = -1
Oxidation 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 = 0
x = -1
Oxidation 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 manganese
MnSO4	+2
MnO2	+4
K2MnO4	+6
KMnO4	+7
Compound	Oxidation number of chromium
K2CrO4	+6
K2Cr2O7	+6

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

Compound	Oxidation number of chlorine
HCl	-1
ClO2	+4
HClO4	+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.

Read More

Redox reaction

Rusting of Iron Nails

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

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.
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.

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 oxidised.

Read More