Reactivity 3
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What are the general rules for assigning oxidation states?
Elements in their ‘elemental state’ are 0. The sum of all individual oxidation states will equal the total charge on the species. Oxygen is usually -2.
What’s unusual about the oxidation states in propane?
The unusual aspect of oxidation states in propane is that the individual carbon atoms have different oxidation states, even though the overall average oxidation state for carbon in the molecule is a non-integer value.Average Oxidation State: The average oxidation state for carbon is 2.67.
Define oxidation and reduction
Oxidation is the loss of electrons by a substance, resulting in an increase in its oxidation state. A common mnemonic is “OIL” (Oxidation Is Loss). Reduction is the gain of electrons by a substance, resulting in a decrease in its oxidation state. A common mnemonic is “RIG” (Reduction Is Gain).
Identify the species being oxidised and reduced in this reaction: Zn (s) + HCl (aq) → ZnCl2 (aq) + H2 (g)
Species oxidised: Zn – it goes from 0 to +2 Species reduced: H⁺ (from HCl) – it goes from +1 to 0
What is an oxidising agent?
An oxidising agent is a substance that oxidises another atom or ion by causing it to lose electrons An oxidising agent itself gets reduced Oxidising agents gain electrons
What is a reducing agent?
A reducing agent is a substance that reduces another atom or ion by causing it to gain electrons A reducing agent itself gets oxidised – loses/donates electrons Therefore, the oxidation number of the reducing agent increases
Deduce the redox half-equations in this reaction: Zn + Cu2+ → Zn2+ + Cu
Oxidation half-equation Zn → Zn²⁺ + 2e⁻ Reduction half-equation Cu²⁺ + 2e⁻ → Cu
What is a redox reaction?
For a reaction to be recognised as a redox reaction, there must be both an oxidising and reducing agent Some substances can act both as oxidising and reducing agents Their nature depends on: What they are reacting with The reaction conditions
The following reaction was carried out in acidic solution: NO3- (aq) + Cu (s) → NO (g) + Cu2+ (aq) Oxidation half-equation: Reduction half equation: Balanced ionic equation:
H2O and H+ will be required to balance half equations that contain 2 or more elements Oxidation: Cu → Cu2+ + 2e- Reduction: NO3- + 4H+ + 3e- → NO + 2H2O Balanced ionic equation: 3Cu + 2NO3- + 8H+ → 3Cu2+ + 2NO + 4H2O
Why oxidation and reduction always occur together
Electrons lost by one species must be gained by another
Purpose of half-equations
To separate oxidation and reduction steps and balance charge and atoms
Steps to balance half-equations in acid
Balance atoms then add H2O then add H+ then add e-
Periodic trend for ease of oxidation in metals
Metals lower and to the left oxidize more easily
Periodic trend for ease of reduction in halogens
Reactivity increases up the group as electronegativity increases
Why reactive metals react with acids
Metals oxidize and H+ is reduced to H2
General redox reaction of metal plus acid
Metal forms metal ions and hydrogen gas is released
Definition of corrosion as redox
Metal is oxidized by oxygen or water to metal ions
Oxidation occurs where in a voltaic cell
At the anode
Reduction occurs where in a voltaic cell
At the cathode
Electron flow direction in a voltaic cell
From anode to cathode externally
Purpose of the salt bridge
Maintains electrical neutrality by allowing ion migration
Definition of primary cell
Non-rechargeable cell where redox reactions are not reversible
Definition of secondary cell
Rechargeable cell where redox reactions can be reversed by electricity
Oxidation occurs where in an electrolytic cell
At the positive electrode