Redox I

Question 1

Oxidation Number

Answer 1

The oxidation number is the charge that the element would have if the compound were fully ionic

• Written as the sign + or - followed by a number
• In non-ionic compounds the bonding electrons are assigned to the more electronegative atom in the bond
• Increase becomes more positive - OXIDISED
• Decrease becomes less positive - REDUCED

Oxidation number rules

• For a molecule the sum of the oxidation numbers of all the atoms must equal zero
• For a complex ion the sum must equal the overall charge on the complex ion
• Oxidation number of fluorine is always -1 as it always gains electrons in bonding
• Oxidation number of oxygen in compounds is always -2 except for when it is bonded to fluorine (+2) or is present as a peroxide (-1) or superoxide.
• Oxidation number of hydrogen is +1 except in metal hydrides (-2)
• Shared pair of electrons is assumed to be under control of more electronegative atom

Question 2

Oxidation

Answer 2

Oxidation is the loss of electrons

When a substance is oxidised, it will lose electrons to another substance

An oxidising agent is a chemical reagent which can oxidise other atoms, molecules or ions by taking electrons away from them

e.g.

• Oxygen
• Chlorine
• Nitric acid
• Potassium manganate
• Potassium dichromate
• Hydrogen peroxide

Question 3

Reduction

Answer 3

Reduction is the gain of electrons

When a substance is reduced, it will gain electrons from another substance

A reducing agent is a chemical reagent which can reduce other atoms, molecules or ions by giving them electrons

e.g.

• Hydrogen
• Sulfur dioxide
• Zinc/iron in acid

When concentrated sulfuric acid is added to sodium chloride, it forms sodium hydrogen sulfate and hydrochloric acid

Chloride isn’t a good enough reducing agent

Question 4

Disproportionation Reaction

Answer 4

Disproportionation is when an element in a single species is both oxidised and reduced simultaneously in the same reaction

When hydrogen peroxide is decomposed..

• Half of the oxygen is reduced from -1 to -2 in water
• The other half is oxidised from the -1 to 0 state in oxygen gas

Reactions of halogens with cold alkali

• With cold alkali, halogens react to form a mixture of the halide and halate (I) salts
• 2OH- + Cl2 → Cl- + ClO- + H2O
• The chlorate salts are good oxidising agents as the chlorine readily accepts electrons to become reduced

Question 5

Thermal Decomposition of Group 1 and 2 Nitrates and Carbonates

Answer 5

1. In order to test any gas given off with lime water, use the apparatus below
2. Begin heating the tube using a gentle blue flame and then increase the temperature of the flame by opening the air hole more
3. Remember to remove the delivery tube from the limewater before the bunsen is removed in order to prevent suck-back of lime
4. water into the hot test tube
5. Other tests such as a glowing splint of damp indicator paper can be carried out simply by heating the solid in a test tube using test tube holders. Strong heat and perseverance is necessary in some cases

• As you go down the group, the carbonates have to be heated more strongly before they will decompose
• As you go down the group, the nitrates have to be heated more strongly before they will decompose

Question 6

Flame Tests for Group 1 and 2 elements

Answer 6

1. Clean a nichrome wire by dipping it into concentrated HCl and then holding it in a hot Bunsen flame. Repeat until the wire doesn’t produce any colour in the flame
2. When the wire is clean, dip it into the HCl and then dip into the solid you are testing
3. Hold the wire in the flame and observe the colour of flame

Lithium (Li+) → Bright Red

Sodium (Na+) → Yellow

Calcium (Ca2+) → Brick red

Potassium (K+) → Lilac

Strontium (Sr2+) → Bright red

Barium (Ba2+) → Green

Magnesium and beryllium do not have flame test colours.