Unit: Electrochemistry Half Reactions and Redox Equations

Oxidation Is Loss of electrons Reduction Gain of electrons Half-Reactions Remember OIL RIG… Oxidation Is Loss of electrons Reduction Gain of electrons

Half-Reaction Example 2HCl + Mg  MgCl2 + H2 1. Label all oxidation numbers 2. Determine what is reduced and oxidized reduction = gain of electrons (more negative #) oxidation = loss of electrons (more positive #)

Half-Reaction Definitions 2HCl + Mg  MgCl2 + H2 Reducing agent: gives electrons and is OXIDIZED Oxidizing agent: removes electrons and is REDUCED

Example: 2HCl + Mg  MgCl2 + H2 The oxidation number of hydrogen becomes more negative, so hydrogen is reduced. Therefore, hydrogen is the oxidizing agent (it oxides another element). The oxidation number of magnesium becomes more positive, so magnesium is oxidized. Therefore, magnesium is the reducing agent (it reduces another element).

Balancing Half Reactions in Acidic Solutions 1. Label all oxidation numbers. 2. Determine which element is reduced and which is oxidized. 3. Write the two half reactions (follow solubility rules!) and do the following to each: Balance element in question using coefficients Balance oxygen by adding H2O Balance hydrogen by adding H+ [leave a blank line here!] Balance charge by adding e- Multiply reactions so that number of e- are equal in both

Balancing Redox Reactions in Acidic Solutions Example #1 H2S + HNO3  H2SO4 + NO2 + H2O 2. H2S  SO4-2 3. H2S + 4H2O  SO4-2 + 10H+

Balancing Redox Reactions in Acidic Solutions 1. Label all oxidation numbers. 2. Determine which element is reduced and which is oxidized. 3. Write the two half reactions (follow solubility rules!) and do the following to each: Balance element in question using coefficients Balance oxygen by adding H2O Balance hydrogen by adding H+ [leave a blank line here!] Balance charge by adding e- Multiply reactions so that number of e- are equal in both 4. Add reactions together, recombine ions, cancel things on both sides, and make the equation look like the original

Balancing Redox Reactions in Acidic Solutions Example #1 H2S + HNO3  H2SO4 + NO2 + H2O

Cr2O7-2 (aq) + Cl- (aq)  Cr+3 (aq) + Cl2 (g) Balancing Redox Reactions in Acidic Solutions Example #2 Cr2O7-2 (aq) + Cl- (aq)  Cr+3 (aq) + Cl2 (g)

H2S + HNO3  H2SO4 + NO2 + H2O 1. Assign oxidation numbers. 2. Write the half-reactions: - If a compound contains an element that changes oxidation #, write that compound in ionic form (follow solubility rules) 3. Balance: - Balance element in question using coefficients - Balance oxygen by adding H2O - Balance hydrogen by adding H+ ions (in basic solutions add OH-) We can add water or H+ because we are assuming we are doing this in an acidic, aqueous solution where these species are readily available 2. H2S  SO4-2 3. H2S + 4H2O  SO4-2 + 10H+

H2S + HNO3  H2SO4 + NO2 + H2O 4. Add electrons to balance the charge 5. Repeat process with the other half-reaction 6. Multiply half reactions so that the number of electrons lost equals the number gained 7. Add the half reactions 8. Re-combine the ions to form the compounds found in the original formula. 4. H2S + 4H2O  SO4-2 + 10H+ + 8e- 5. e- + NO3- + 2H+ NO2 + H2O 6. 1(H2S + 4H2O  SO4-2 + 10H+ + 8e-) 8(e- + NO3- + 2H+ NO2 + H2O) 7. 8e- + 8NO3- + 16H+ + H2S + 4H2O  8NO2 + 8H2O + SO4-2 + 10H+ + 8e- After canceling… 8NO3- + H2S + 6H+ 8NO2 + 4H2O + SO4-2 8. LEFT: 8HNO3 + H2S  8NO2 + 4H2O + SO4-2 + 2H+ (add 2H+ to both sides to make 8HNO3) RIGHT: 8HNO3 + H2S  8NO2 + 4H2O + H2SO4

Balancing Redox Equations – Practice #1 Balance the redox equation below: NaBr(aq) + Cl2(g)→ NaCl(aq) + Br2(l)

Balancing Redox Equations – Practice #2 Balance the redox equation below: Fe2O3(s) + CO(g) → Fe(s) + CO2(g)

Balancing Redox Equations – Practice #3 Balance the redox equation below: I2(s) + OCl-(aq) → IO3- (aq) + Cl-(aq)

Balancing Redox Equations – Practice #4 Balance the redox equation below: HCOOH(aq) + MnO4-(aq) → CO2 (g) + Mn2+(aq) Note: HCOOH is a weak acid.

Balancing Redox Reactions in Basic Solutions 1. Label all oxidation numbers. 2. Determine which element is reduced and which is oxidized. 3. Write the two half reactions (follow solubility rules!) and do the following to each: Balance element in question using coefficients Balance oxygen by adding H2O Balance hydrogen by adding H+ Add OH- to neutralize H+ Balance charge by adding e- Multiply reactions so that number of e- are equal in both 4. Add reactions together, recombine ions, cancel things on both sides, and make the equation look like the original

Cr+3 (aq) + O2 (g)  Cr2O7-2 (aq) + H2O2 (l) Balancing Redox Reactions in Basic Solutions Example #1 Cr+3 (aq) + O2 (g)  Cr2O7-2 (aq) + H2O2 (l)