Redox

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• Reduction and oxidation

Note:

Read questions carefully. Answer the entire question. Write what they tell you to write, not what you want to write

• Reduction: gain of electrons from chemical species. The species being reduced is the oxidising agent
  • ${\text{Cl}}_2+2{\text{e}}^{-}+2{\text{Cl}}^{-}$
• Oxidation: loss of electrons from chemical species. The species being oxidised is the reducing agent
  • $2\text{Na}\to 2{\text{Na}}^{+}+2{\text{e}}^{-}$
• Redox: exchange of electrons
• Due to conservation of matter, reduction and oxidation is the exchange of electrons and happens simultaneously
• OILRIG
  • Oxidation
  • Is
  • Loss
  • Reduction
  • Is
  • Gain
• Oxidation number (oxidation state): A number assigned to the atom to show how many electrons are lost, gained or shared unequally (to form a bond)
  • Increase in oxidation number is loss of electrons
  • Decrease in oxidation number is gain in electrons
  • Think charge in the ion

Elements with fixed oxidation states in their compounds	Oxidation state
Hydrogen (except in metal hydrides)	+1 (-1)
Oxygen (except in peroxides)	-2 (-1)
Group I elements	+1
Group II elements	+2
Aluminium	+3
Halide	-1

Oxidation Rules

• For uncombined elements, the oxidation state is zero
• For atoms in elements (when element is combined to itself), the oxidation state is zero
• For simple ions, the oxidation state is the charge on the ion
• For compounds, the sum of the oxidation states (of all present species) is 0
• For more complex ions, the sum of oxidation numbers is the charge on the ion
• Oxidation states can be assigned to elements in covalent compounds even though complete transfer of electrons does not occur

Oxidation number

Put the oxidation number under element, and the total above element, when calculating oxidation states

BE SPECIFIC AND PARTICULAR. DO NOT BE VAGUE. EXPLAIN WELL AND LOGICALLY, E.G. "IT" IS NOT GOOD ENOUGH

Redox titrations are not on the syllabus, but there could be an application of it on WACE papers

Balancing Redox equations in acidic conditions

0. Assign oxidation numbers
1. Write the skeletons of the oxidation and reduction half reactions
2. Balance all elements other than H and O
3. Balance the oxygen atoms by adding $H_{2}O$, where needed
4. Balance the hydrogen atoms by adding $H^{+}$ ions, where needed
5. Balance the charge by adding electrons, $e^{-}$
6. If the number of electrons lost in the oxidation half-reaction is not equal to the number of electrons gained in the reduction half reaction, multiply one or both of the half reactions by a number that will make the number of electrons gained equal to the number of electrons lost
7. Add the 2 half-reactions as if they were mathematical equations. The electrons will always cancel. if the same formulas are found on the opposite sides of the half reactions, you can cancel them. If the same formulas are found on the same side of both half reactions, combine them
8. Check to make sure that the atoms and the charges balance

Displacement reactions of Halogens

Halogens (oxidising agents)	Halides (reducing agents)
$C{l}_2+2\overline{e}\rightleftharpoons$	$2C{l}^{-}$
$B{r}_2+2\overline{e}\rightleftharpoons$	$2B{r}^{-}$
$I_2+2\overline{e}\rightleftharpoons$	$2I^{-}$

If a substance is coloured, it is not due to the halides, hence only halogens and colours are listed on page 5 of the data book

1. $C{l}_2+2B{r}^{-}\to 2C{l}^{-}+B{r}_2$
   1. Ox: $2B{r}^{-}\to B{r}_2+2\overline{e}$
   2. Red: $C{l}_2+2\overline{e}\to 2C{l}^{-}$
2. $C{l}_{2(aq)}+2I^{-}\to 2C{l}^{-}+I_2$
3. $B{r}_{2(aq)}+C{l}^{-}\to \text{no reaction}$
4. $B{r}_{2(aq)}+2I^{-}\to 2B{r}^{-}+I_2$
5. $I_2+C{l}^{-}\to \text{no reaction}$
6. $I_2+B{r}^{-}\to \text{no reaction}$

$E_{\text{cell}}=E_{\text{red}}-E_{\text{ox}}$

• If it is positive, the reaction is feasible