3 - Redox Ⅰ
● ***Reduction: gain of electrons/hydrogen (in chemical reaction); decrease in
oxidation number
● ***Oxidation: loss of electrons/hydrogen (```); increase ```
○ An example of continuous oxidation: NH3 ⇌ N2 ⇌ NO2- ⇌ NO3-
-3 0 +3 +5
○ Oxidation of oxygen eg O2 + 2F2 → 2F2O
0 0 -1+2
● ***Oxidation number: theoretical charge of an atom in compound (if compound was
purely ionic)
○ Ie elements have oxidation number of 0
○ Sum of all oxidation states of all atoms in molecule/ion = overall charge of
molecule/ion
○ RULES:
1. Hydrogen = +1 (otherwise hydride = -1)
2. Oxygen = -2 (unless peroxide [O22-] = -1, superoxide [O2-] = -½; others
+2)
3. Group 7 = -1
4. Group 1 = +1, Group 2 = +2
○ No change in oxidation number (in reaction) - spectator ion
■ No change FOR ALL IONS - NOT redox reaction
● ***Oxidising agent: oxidises sth else during reaction (while itself being reduced)
○ Ie strong oxidising agents are more electronegative (stronger tendency to
gain electrons)
● ***Reducing agent: reduces ```` (``` oxidised)
○ Ie strong reducing agents are less electronegative (stronger tendency to lose
electrons)
● ***Disproportionation: when a species is simultaneously oxidised and reduced at
the same time
○ Eg 2Cu+ (aq) → Cu (s) + Cu2+ (aq) {copper (I) unstable in solution}
○ Eg2 3Fe2+ (aq) → 2Fe3+ + Fe (s)
○ Eg3 Cl2 + 2NaOH → NaClO + NaCl + H2O {in cold, aqueous NaOH}
■ Eg3b 3Cl2 + 6NaOH → 5NaClO3 + NaCl + 3H2O {in hot, concd NaOH}
○ Eg4 2H2O2 → 2H2O + O2 {decomposition of hydrogen peroxide}
○ Eg5 Cl2 + H2O → HCl + HClO {production of commercial bleach}
Half equations → ionic (full) equations
● Eg oxidation half equation: Zn (s) → Zn2+ (aq) + 2e-
Reduction ``: Cu2+ (aq) + 2e- → Cu (s)
⇒ Sum to get redox equation (e- eliminated): Cu2+ (aq) + Zn (s) → Cu (s) + Zn2+ (aq)
{Equations can ONLY be combined directly when e- are on opposite sides of
equations}
● RULES for balancing ionic equations:
1. (Balance chemical species - in terms of quantity)
1
● ***Reduction: gain of electrons/hydrogen (in chemical reaction); decrease in
oxidation number
● ***Oxidation: loss of electrons/hydrogen (```); increase ```
○ An example of continuous oxidation: NH3 ⇌ N2 ⇌ NO2- ⇌ NO3-
-3 0 +3 +5
○ Oxidation of oxygen eg O2 + 2F2 → 2F2O
0 0 -1+2
● ***Oxidation number: theoretical charge of an atom in compound (if compound was
purely ionic)
○ Ie elements have oxidation number of 0
○ Sum of all oxidation states of all atoms in molecule/ion = overall charge of
molecule/ion
○ RULES:
1. Hydrogen = +1 (otherwise hydride = -1)
2. Oxygen = -2 (unless peroxide [O22-] = -1, superoxide [O2-] = -½; others
+2)
3. Group 7 = -1
4. Group 1 = +1, Group 2 = +2
○ No change in oxidation number (in reaction) - spectator ion
■ No change FOR ALL IONS - NOT redox reaction
● ***Oxidising agent: oxidises sth else during reaction (while itself being reduced)
○ Ie strong oxidising agents are more electronegative (stronger tendency to
gain electrons)
● ***Reducing agent: reduces ```` (``` oxidised)
○ Ie strong reducing agents are less electronegative (stronger tendency to lose
electrons)
● ***Disproportionation: when a species is simultaneously oxidised and reduced at
the same time
○ Eg 2Cu+ (aq) → Cu (s) + Cu2+ (aq) {copper (I) unstable in solution}
○ Eg2 3Fe2+ (aq) → 2Fe3+ + Fe (s)
○ Eg3 Cl2 + 2NaOH → NaClO + NaCl + H2O {in cold, aqueous NaOH}
■ Eg3b 3Cl2 + 6NaOH → 5NaClO3 + NaCl + 3H2O {in hot, concd NaOH}
○ Eg4 2H2O2 → 2H2O + O2 {decomposition of hydrogen peroxide}
○ Eg5 Cl2 + H2O → HCl + HClO {production of commercial bleach}
Half equations → ionic (full) equations
● Eg oxidation half equation: Zn (s) → Zn2+ (aq) + 2e-
Reduction ``: Cu2+ (aq) + 2e- → Cu (s)
⇒ Sum to get redox equation (e- eliminated): Cu2+ (aq) + Zn (s) → Cu (s) + Zn2+ (aq)
{Equations can ONLY be combined directly when e- are on opposite sides of
equations}
● RULES for balancing ionic equations:
1. (Balance chemical species - in terms of quantity)
1
