Redox reactiosn

  • A redox reaction is a process in which electrons are transferred from one species to another. 

  • The species that donates the electrons is said to be oxidised, and is called a reducing agent as it causes another species to be reduced.

  • The species that gains electrons is said to be reduced, and is called a oxidising agent as it causes oxidation of another species. An easy way to remember this is OILRIG.

  • Oxidation Is Loss Reduction Is Gain (OIL RIG)

 

Note that oxidation and reduction always occur simultaneously.

  • Reduction

    • Loss of oxygen

    • Gain of hydrogen

    • Gain of electrons

    • Decrease in oxidation number

  • Oxidation

    • Gain of oxygen

    • Loss of hydrogen

    • Loss of electrons

    • Increase in oxidation number

 

Different types of redox reactions

  • Metal-metal ion displacement ⇒ electrons transfer from more reactive metal to less reactive metal ions (more reactive metal is oxidised, less reactive reduced)

  • Halogen-halide ion displacement ⇒ halogen (group 17) becomes reduced, halide ions of less reactive halogen oxidised

    • It is more convenient to use halogen reagents in aqueous solution form

  • Combustion ⇒ oxidisation of fuel and reduction of oxygen gas

  • Corrosion ⇒ Metal is oxidised, and oxygen gas is reduced

    • More reactive metals have greater tendency to corrode

    • Some metals naturally form a thin protective oxide coating which forms when exposed to air, hence protecting underlying metal from corrosion

 

 

 

Half Equations

  • The processes of oxidation and reduction can be shown through half equations

  • Used to determine whether a species is oxidised or reduced

  • Each half describes the oxidation process -gaining electrons - and the reduction process -gaining electrons.

  • If electrons are included in the reactants it is reduction

  • If electrons are included in the products it is oxidation

  • Each half can be added to give the overall net equation

For example

  • A:     +     B     →     B: -2      +     A +2

  • Oxidation Half:     A:    →   A +2    +    2e-

  • Reduction Half:    B   +   2e-    →   B: -2

 

 

 

 

 

 

 

 

Oxidation Numbers

  • Oxidation numbers are numerical values assigned to atoms, allowing us to keep track of electron exchange, and allow us to determine whether or not a redox reaction is taking place.

  • There are 5 rules when it comes to oxidation numbers (they are imaginary numbers)

 

  1. All substances in their elemental state get a oxidation number of 0  (how they are found in nature). For examples things like metals in their natural form. Mg(s) = 0 H2 = 0

 

  1. Monoatomic ions = charge of their ion. Fe3+ = 3   02- = -2

 

  1. The oxidation numbers for combined Oxygen = -2  (peroxides where it equals +2)

 

  1. Combined Hydrogen = +1 (metal hydrides is - 1)

 

  1. If you have a polyatomic ion, and you sum up all oxidation numbers = charge of the ion ex. S04 2-

Ox(s) + four lots of -2 = -2

 

Note if oxidation number is increasing it is becoming oxidised, if it is decreasing it is reducing (reduction).

 

Finding full equations

  1. Oxidation numbers

  2. OIL RIG or is oxidation number increasing (ox) or decreasing (red)

  3. Balance only the atoms that change oxidation numbers

  4. Only Add H20 to balance the oxygen atoms only  (if there is a presence of oxygen atoms)

  5. Only Add H+ (Acidic conditions) to balance hydrogen atoms(if there is a presence of hydrogen atoms)

  6. Add electrons to balance overall charge

 

 

Y(s) + -> X(s) + Y 3+

 

 

Oxidising agents

  • If an oxidising agent/oxidant causes oxidation to occur in another substance, itself must be reduced

  • Vice versa reducing agent/ reductant causes reduction to occur in another substance, and itself must be oxidised.