Balancing Redox Equations: following the electrons

1. Balancing Redox Equations: following the electrons Review: Oxidation and reduction Oxidation numbers

2. Review: Oxidation - reduction • Oxidation is loss of electrons • Reduction is gain of electrons • Oxidation is always accompanied by reduction • The total number of electrons is kept constant • Oxidizing agents oxidize and are themselves reduced • Reducing agents reduce and are themselves oxidized

3. Nuggets of redox processes • Where there is oxidation there is always reduction

4. Oxidation numbers review • Metals are more 'cation-like' • Have positive oxidation numbers • Nonmetals are 'anion-like' • Have negative oxidation numbers. • Oxidation number is the number of electrons gained or lost by the element in making a compound

5. Predicting oxidation numbers • Oxidation number of atoms in element is zero in all cases • Oxidation number of element in monatomic ion is equal to the charge • sum of the oxidation numbers in a compound is zero • sum of oxidation numbers in polyatomic ion is equal to the charge • F has oxidation number –1 • H has oxidn no. +1; except in metal hydrides where it is –1 • Oxygen is usually –2. Except: • O is –1 in hydrogen peroxide, and other peroxides • O is –1/2 in superoxides KO2 • In OF2 O is +2

6. Position of element in periodic table determines oxidation number • G1A is +1 • G2A is +2 • G3A is +3 (some rare exceptions) • G5A are –3 in compounds with metals, H or with NH4+. Exceptions are in compounds to the right; in which case use rules 3 and 4. • G6A below O are –2 in binary compounds with metals, H or NH4+. When they are combined with O or with a lighter halogen, use rules 3 and 4. • G7A elements are –1 in binary compounds with metals, H or NH4+ or with a heavier halogen. When combined with O or a lighter halogen, use rules 3 and 4.

7. Redox equations • Net ionic equations summarize the essentials of a reaction without including all the particles present • Redox equations are a subset which involve electron transfer • Without being given all the information, balancing redox equations involves balancing electron flow

8. Balancing redox equations: systematic methods • Oxidation number method – tracking changes in the oxidation numbers • Half-reaction method – tracking changes in the flow of electrons • Same principles, different emphasis • We will examine the half-reaction method

9. The Half-Reaction method • Any redox process can be written as the sum of two half reactions: one for the oxidation and one for the reduction

10. Six habits of the redox equation balancer

11. STEP 1: the unbalanced equation • Dichromate ion reacts with chloride ion to produce chlorine and chromium (III)

12. STEP 2: identify the oxidized and reduced and write the half reactions • Oxidation half-reaction • Reduction half-reaction

13. STEP 3: Balance the half reactions • Oxidation • Reduction

14. Material balance with H2O and H+ or OH- • Strategy: add H2O to the side that lacks for O and add H+ (the reaction is in acid solution) to the other side • In basic solution we add OH- and H2O instead of H2O and H+ respectively • Test equation for both atoms and charges

15. STEP 4: Material balance • Add H2O to the side lacking O and add H+ to the other side (for reactions in acid solution) • Oxidation reaction – unchanged • Reduction reaction

16. STEP 5: Balance half-reactions for charge by addition of electrons • Balance charges on both sides of each half-reaction 2 x -1 = 2 x -1 14 x +1 + -2 + 6 x -1 = 2 x +3

17. STEP 5 cont: Multiply by factors to balance total electrons • Overall change in electrons must be zero • Multiply the oxidation half reaction by 3 3 x 2 = 6

18. STEP 6: Add half reactions and eliminate common items + = Electrons cancel both sides Atoms and charges balance

19. Balanced molecular equation • Add in the spectators: there will always be space. Reagents were K2Cr2O7, NaCl and H2SO4 • Net ionic equation • Balanced molecular equation