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 constantThe 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 Oxidizing agent Reducing agent Is itself reduced Is itself oxidized Gains electrons Loses electrons Causes oxidation Causes 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 KO 2  In OF 2 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 NH 4+. 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 NH 4+. 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 NH 4+ 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 H 2 O and H + or OH -  Strategy: add H 2 O 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 H 2 O instead of H 2 O and H + respectively  Test equation for both atoms and charges

15. STEP 4: Material balance  Add H 2 O 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 K 2 Cr 2 O 7, NaCl and H 2 SO 4  Net ionic equation  Balanced molecular equation