• 1.5 Oxidation and• Reduction

Learning Outcomes

• Introduction to oxidation and reduction: simple examples only, e.g. Na with Cl2, Mg with O2, Zn with Cu2+. Oxidation and reduction in terms of loss and gain of electrons.

• Oxidising and reducing agents.• The electrochemical series as a series of metals arranged in order of

their ability to be oxidised (reactions, other than displacement reactions, not required).

• Electrolysis of (i) copper sulfate solution with copper electrodes and (ii) acidified water with inert electrodes. (Half equations only required.)

Oxidation and reduction

• Oxidation = addition of oxygen to a substance

• C + O2 CO2

• Reduction is loss of oxygen or addition of hydrogen

• CuO + H2 Cu + H2O

examples

• Sodium + chlorine sodium chloride

• Na + Cl Na+ + Cl-

• Na loses an electron [oxidised]

• Cl gains an electron [reduced]

Example 2

• Magnesium +oxygen magnesium oxide

• Mg + O MgO

• Mg Mg+2 loses 2 electrons [oxidation]• O O-2 gains 2 electrons [reduction]

Example 3

• Zinc +copper sulphate Zinc sulphate+ Copper

• Zn + Cu+2 Zn+2 + Cu

• Zinc loses electrons (oxidised)

• Copper gains electrons (reduced)

Oxidising agent

• A substance that causes oxidation in another substance

Reducing agent• A substance that causes reduction in another

substance.

• Oxidation is loss of electrons; • Reduction is gain of electrons• CuO + H2 Cu + H2O

• CuO Cu+2 and O-2

• Cu+2 Cu [gains 2 electrons] reduced• H2 H2

+2[loses 2 electrons] oxidised

• O-2 O-2 [ no change]

Oxidation numbers

• The charge that an atom has or appears to have assuming that the compound is ionic.

• Electrons always go the the most electronegative element

Oxidation number rules 1

• Elements on their own = 0• H2 = 0

• Zn = 0• Cl2 = 0

Oxidation number rules 2

• Ions = same as charge• Cu +2 = +2• O-2 = -2• Cl-1 = -1

Oxidation number rules 3

• Charges of all elements in a compound = 0• CuSO4

• Cu = +2• S = +6• O4 = -8 [O = -2]

• Total = +2 +6 –8 = 0

Oxidation number rules 4

• Oxygen = -2• Exceptions are• peroxides O = -1 [H2O2, Na2O2 ]

• OF2 O = +2, F = -1

Oxidation number rules 5

• Hydrogen = +1• Exceptions are the metal hydrides• NaH Na = +1, H = -1

Oxidation number rules 6

• Halogens [ Cl, F, I, Br] are always –1 except when joined to more electropositice element

• Cl2O

• Cl = +1, O = -2

Oxidation number rules 7

• In a complex ion the sum of all the charges = the chartge on the ion.

• SO4-2

• S = +6, O4 = -8 [O = -2]

• +6 –8 = -2

redox

• Oxidation is an increase on oxidation number

• Reduction is a decrease in oxidation number.

Electrochemical Series

• Electrochemical Series – Elements listed in order of ability to be oxidised

Metals• King [K] • Neptune [Na] • Caught [Ca] • Many [Mg] • Angry [Al] • Zulus [Zn] • Fighting [Fe] • Police [Pb] • Constables [Cu] • Having [Hg] • Asthma [Ag] • Attacks [Au]

Metals above hydrogen in the

Reactivity Series react with acids to produce hydrogen

gas.

Zinc

Potassium

Sodium

Displacement of metals

• Displacement reactions occur when a metal from the electrochemical series is mixed with the ions of a metal lower down in the series. The atoms of the more reactive metal push their electrons on to ions of the less reactive metal.

Displacement

• More reactive metal displaces less reactive from a solution

• Mg + CuSO4 = MgSO4 + Cu

• Mg + Cu+2 Mg+2 + Cu• Mg loses electrons (Oxidised)• Cu+2 gains electrons (reduced)

Learning Outcomes

• Rusting of iron.• Swimming-pool water treatment.• Use of scrap iron to extract copper.• Electroplating. Purification of copper.• Chrome and nickel plating. Cutlery.

Rust

• Rust is the formation of iron oxides (usually red oxides), formed by the reaction of iron and oxygen in the presence of water or air moisture.

• Oxidation

Swimming pools

• The water in swimming pools is kept sterile by the addition of oxidizing agents, chlorine or chlorine compounds, which kill microorganisms by oxidation. The active agent is usually chloric(1) acid (HOCl). It may be formed in two ways

• 1. Direct chlorination of the water: • Cl 2(aq) +H2O (l) HOCl (aq) + Cl−

(aq) + H+ (aq)

• Note that when the Cl2 reacts with the water it is both

oxidized and reduced

Swimming pools

• 2. The addition of sodium chlorate(I) [sodium hypochlorite]:

• NaOCl (s) + H2O (l) Na+ (aq) + OH−

(aq) + HOCl (aq)

• Nowadays chlorine is not used, mainly on grounds of safety. Pools are sterilized with chlorine compounds, which produce chloric(I) acid when they dissolve in water. These compounds act in essentially the same way as chlorine. Sodium chlorate(I) is one such compound.

Use of scrap iron to extract copper.

• (Dissolved CuSO4) + (Metallic Fe) ==> (Dissolved FeSO4) + (Metallic Cu)

Electrolysis

• Chemical reaction caused by the passage of an electric current through a liquid known as the electrolyte

Definitions• Electrolyte - liquid in which electrolysis takes place. Usually an ionic solution but it can also be a

fused [melted] ionic compound

• Anode - positive electrode. Positive because the battery sucks electrons out of it•  • Cathode. Negative electrode. Negative because the battery pumps electrons into it.•  • Anion - negative ion. Called anion because it is attracted to the opposite charge of the anode

• Cation - positive ion. Called cation because it is attracted to the opposite charge of the cathode.

• Inert Electrodes - do not react with the electrolyte Graphite and Pt

• Active electrodes - react with electrolyte e.g. Copper and iron

 

Electrolysis

Electroplating

• Electroplating• Covering cathode in metal e.g. Cu by

making it cathode in copper sulphate solution

Copper plating

Copper Plating

Anode reaction

• Cu(s) = Cu2+(aq) + 2e-

• Anode loses mass as copper dissolves off• Impurities [Au, Ag, Pt etc.] fall to bottom

Cathode reaction

• Cu2+(aq) + 2e- = Cu(s)

• Cathode gains mass as Cu is deposited on it• Cu is 99.9% pure

 

Learning Outcomes

• Mandatory experiment 1.2 (half equations only required, e.g. 2Br– – 2e– → Br2).

• Demonstration of ionic movement.• Demonstration of electrolysis of aqueous sodium sulfate

(using universal indicator) • and of aqueous potassium iodide (using phenolphthalein

indicator) with inert electrodes. (Half equations only required.)

Ionic Movement• During electrolysis of a solution of Copper Chromate in

dil. Hydrochloric acid, positive ions (cations) are attracted to the negative electrode (cathode) and negative ions (anions) are attracted to the positive electrode (anode). If these ions are coloured, their movement may be observed visually.

• Examples of coloured ions include; • copper(II) [Cu2+] - blue • chromate(VI) [CrO42- ] – yellow

Q & A to Ionic Movement Expt

• (1) What colour is the copper(II) chromate solution? • Copper(II) chromate solution is an olive green colour.

• (2) What colour is observed at the positive electrode after the power supply has been turned on for some time? • A yellow colour is observed at the positive electrode.

• (3) What colour is observed at the negative electrode after the power supply has been turned on for some time? • A blue colour is observed at the negative electrode.

• (4) Explain in terms of the movement of ions why different colours are formed at each electrode. • When the circuit is completed, positive copper ions (Cu2+) are attracted to the negative electrode. These ions have a

blue colour. Similarly negative chromate(VI) ions (CrO42-) are attracted to the positive electrode. These ions are coloured orange.

• (5) What is the function of the dilute hydrochloric acid? • The dilute hydrochloric acid is required to complete the circuit.

Electrolysis of Sodium sulphate

• Solution of Na2SO4 + universal indicator

• H+ ions are produced at the positive electrode (oxidation of O2- in water) while OH- ions are produced at the negative electrode as the H+ in water is reduced to H2(g).

Sodium Sulphate and Universal Indicator

Electrolysis of Sodium Sulphate

• Red is acid at the positive electrode• 2H2O(l) O2(g) + 4H+(aq) + 4 e-

• lose electrons = oxidation = anode•  • Purple is base at the negative electrode• H2O(l) + 2 e- H2(g) + 2OH-(aq)

• gain electrons = reduction = cathode

Electrolysis of Potassium Iodide

• Solution of KI + phenolphthalein

• Brown I2(s) forms at the positive electrode and some yellow/orange I3

- forms in solution. At the negative electrode, H+ is again reduced to H2(g) and the phenolphthalein turns pink due to the OH- ions.

Electrolysis of Potassium Iodide

Electrolysis of Potassium Iodide

• KI K+ + I-• Iodide loses electrons Brown iodine• 2I- I2 + 2e- Anode, Oxidation

• H2O H+ + OH-

• OH- is basic , Phenolphthalein Purple.