Red uction – Ox idation Reactions “REDOX”

1. Reduction – Oxidation Reactions “REDOX” • Examples of Redox Reactions: • Formation, decomposition, combustion, single replacement, cellular respiration, photosynthesis, (NOT double replacement)

2. Predicting Redox Reactions 13.2 • REMEMBER a redox reaction is a transfer of valence electrons from one substance to another. • Chemists use the analogy of a tug-of-war to describe the competition for electrons • In a reaction, each entity pulls on the same electrons • If one entity is able to pull electrons away from the other, a spontaneous reaction occurs • If one entity is unable to pull electrons away from the other, no reaction occurs (non-spontaneous)

3. How do we know…? • Without mixing all possible reactants and observing any evidence of reaction, how can we predict if a reaction will occur? (spontaneous or nonspontaneous)?? • If a reaction occurs, what will be produced?? • The answers to these questions cannot be obtained easily using redox theory. • By observing many successful and unsuccessful reactions, patterns emerge and empirical generalizations can be made. • But first before looking at patterns we need to understand some terms: OA (oxidizing agent) and RA (reducing agent)…

4. Oxidizing and Reducing Agents • In any redox reaction, an electron transfer occurs: • One reactant is oxidized, and one reactant is reduced.

5. Terms to identify e- transfer. • Rather than saying “the reactant that is oxidized” and “the reactant that is reduced,” chemists use the terms reducing agent (RA) and oxidizing agent (OA) Note: Remember when an atom becomes more + it is losing electrons.

6. A reducing agent causes reduction by donating (losing) electrons. During this process it is oxidized. • An oxidizing agent causes oxidation by removing (gaining) electrons. During this process it is reduced. • It is important to note that oxidation and reduction are processes, and oxidizing agents and reducing agents are substances.

7. Note: From ancient times “reduced” meant to take a substance (compound) and “reduce” it to its pure form. Reduced = is gaining electrons. Also anciently the term “oxidizing agent” was thought to only oxygen but then found that others (eg. Halogens) can also oxidize/corrode metals.

8. Redox Terms • Silver ions were reduced to silver metal by reaction with copper metal. Simultaneously, copper metal was oxidized to copper(II) ions by reaction with silver ions. • If Ag+(aq) is reduced it is the: • If Cu(s) is oxidized it is the: OXIDIZING AGENT (OA) REDUCING AGENT (RA) It is important to note that oxidation and reduction are processes, and oxidizing agents and reducing agents are substances.

9. REDOX Reactions … so far Reduction Oxidation Historically, the formation of a metal from its “ore” (or oxide) I.e. nickel(II) oxide is reduced by hydrogen gas to nickel metal NiO(s) + H2(g) Ni(s) + H2O(l) Ni +2  Nio A gainof electrons occurs (so the entity becomes more negative) Electrons are shown as the reactantin the half-reaction A species undergoing reduction will be responsible for the oxidation of another entity – and is therefore classified as an oxidizing agent (OA) • Historically, reactions with oxygen • I.e. iron reacts with oxygen to produce iron(III) oxide 4 Fe(s) + O2(g) Fe2O3(s) Fe 0  Fe+3 • A lossof electrons occurs (so the entity becomes more positive) • Electrons are shown as the productin the half-reaction • A species undergoing oxidation will be responsible for the reduction of another entity – and is therefore classified as an reducing agent (RA)

10. Redox Terms • Summary so far: • The substance that is reduced (gains electrons) is also known as the oxidizing agent • The substance that is oxidized (loses electrons) is also knows as the reducing agent (Is Reduced) (Is Oxidized) • Question: If a substance is a very strong oxidizing agent, what • does this mean in terms of electrons? • The substance has a very strong attraction for electrons. • Question: If a substance is a very strong reducing agent, what • does this mean in terms of electrons? • The substance has a weak attraction for its electrons, which are easily removed

11. Development of a Redox Table • In the past we have generally assumed that all single replacement reactions are spontaneous. But now we can’t assume… so… • How do you know when a chemical reaction will occur spontaneously without actually doing the reaction? • Experiments and a redox table.

12. Redox Table • A reaction is considered spontaneous if it occurs on its own • A reduction ½ reaction table is useful in predicting the spontaneity of a reaction • Reduction Tables show reduction ½ reactions in the forward direction, therefore all the reactants will be oxidizing agents • If we list the OA’s from (P.569 look at table 1) in decreasing order of strength, we create a reduction ½ reaction table: Ag+(aq) + 1 e-Ag(s) Cu2+(aq) + 2 e-Cu(s) Pb2+(aq) + 2 e-Pb(s) Zn2+(aq) + 2 e-Zn(s) WRA SOA SRA WOA

13. Strongest oxidizing agent at the top left • Strongest reducing agent at the bottom right • For metal ions (oxidizing agents), the half-reactions are read from left to right in the table. • For metal atoms (reducing agents), the half-reactions are read from right to left.

14. Reading a Redox Table • The table lists the relative strengths of oxidizing and reducing agents. • Listed as reductions (from left to right) in the form: • The strongest oxidizing agent is listed at the top left • The strongest reducing agent is listed at the bottom right

15. Summary • An oxidizing agent causes oxidation by removing (gaining) electrons from substance in a redox reaction. In the process the oxidizing agent is reduced. • A reducing reagent promotes reduction by donating (loosing) electrons to another substance in a redox reaction. In the process the reducing agent is oxidized. • A table of relative strengths of oxidizing and reducing agents –known as a redox table – is, by convention, listed as reductions (from left to right) in the form , with the strongest oxidizing agent at the top left and strongest reducing agent at the bottom right of the table.

16. Homework… • p.571 #1-3,6,7

17. The redox spontaneity rule states that a spontaneous redox reaction occurs only if the oxidizing agent is above the reducing agent in the redox table.

18. Thankfully for us… • Chemists have analyzed, experimented, and collected enough data to be able to produce an extended redox table of oxidizing and reducing agents…. And that table is found in the back of your text book (Appendix 1 p. 828 or in Data booklets). • You can use this table as a reference when trying to predict spontaneous redox reactions.

19. Appendix 1 p. 828 or in Data booklets • The relative position of a pair of oxidizing agents and reducing agents indicates whether a reaction will be spontaneous. This picture is from your data booklet reduction ½ reaction table

20. Building Redox Tables #2 • Example 2: Use the following information to create a table of reduction ½ reactions 3 Co 2+(aq) + 2 In(s) 2 In 3+(aq) + 3Co(s) Cu 2+(aq) + Co(s)  Co 2+(aq) + Cu(s) Cu 2+(aq) + Pd(s) no reaction Pd2+(aq) + 2 e-Pd(s) Cu2+(aq) + 2 e-Cu(s) Co2+(aq) + 2 e-Co(s) In3+(aq) + 3 e-In(s) RA OA OA RA RA OA SOA SRA

21. Building Redox Tables #3 • Example 3: Use the following information to create a table of reduction ½ reactions 2 A 3+(aq) + 3 D(s) 3 D2+(aq) + 2A(s) G +(aq) + D(s)no reaction 3 D 2+(aq) + 2 E(s)3 D(s) + 2 E3+(aq) G +(aq) + E(s)no reaction A3+(aq) + 3 e-A(s) D2+(aq) + 2 e-D(s) E3+(aq) + 3 e-E(s) G+(aq) + 1 e-G(s) OA RA A3+ OA RA D2+(aq) D(s) E(s) OA RA G+ OA RA SOA SRA

22. Building Redox Tables • So far we have been using examples where the oxidizing agents are metal ions and the reducing agents are metal atoms. What else could gain or lose electrons? • Non-metal atoms I.e. Cl2(g) + 2e- 2 Cl-(aq) (Cl2(g) could act as a Reducing Agent) • Non-metal ions I.e. 2 Br-(aq) Br2(l) + 2 e- (2Br-(aq) could act as an Oxidizing Agent) • Redox Table Trend • OA’s tend to be metal ions and non-metal atoms • RA’s tend to be metal atoms and non-metal ions • Also, are there any entities that could act as both OA or RA? • Multivalent metals

24. Try Pg. 573 #14

25. Pg. 573 #14 • Example 4: Use the following information to create a table of reduction ½ reactions Ag(s) + Br2(l) AgBr(s) Ag(s) + I2(s)no evidence of reaction Cu2+(aq) + I-(aq)no redox reaction Br2(l) + Cl-(aq)no evidence of reaction Cl2(g) + 2 e- 2Cl-(aq) Br2(l) + 2 e- 2Br-(aq) Ag+(aq) + 1 e-Ag(s) I2(s) + 2 e- 2I-(aq) Cu2+(aq) + 2 e-Cu(s) RA OA Cl- Br2(l) OA RA Ag(s) I2(s) I-(aq) OA RA Cu2+(aq) RA OA SOA SRA

26. Homework…. • Textbook: • p.571 #1-3,6,7 • p.573 #12 • p.574 #15,17,20 • Diploma Questions: • Ex. MC P.25 #35, 37 • Ex. MC P. 27 #39 • Ex. MC P. 29 #42 • MC 2013 #10, 12 • NR 2013 #5 • MC 2012 #12

27. Building Redox Tables • So far we have been using examples where the oxidizing agents are metal ions and the reducing agents are metal atoms. What else could gain or lose electrons? • Non-metal atoms I.e. Cl2(g) + 2e- 2 Cl-(aq) (Cl2(g) could act as a Reducing Agent) • Non-metal ions I.e. 2 Br-(aq) Br2(l) + 2 e- (2Br-(aq) could act as an Oxidizing Agent) • Redox Table Trend • OA’s tend to be metal ions and non-metal atoms • RA’s tend to be metal atoms and non-metal ions • Also, are there any entities that could act as both OA or RA? • Multivalent metals

28. Try Pg. 573 #14

29. Pg. 573 #14 • Example 4: Use the following information to create a table of reduction ½ reactions Ag(s) + Br2(l) AgBr(s) Ag(s) + I2(s)no evidence of reaction Cu2+(aq) + I-(aq)no redox reaction Br2(l) + Cl-(aq)no evidence of reaction Cl2(g) + 2 e- 2Cl-(aq) Br2(l) + 2 e- 2Br-(aq) Ag+(aq) + 1 e-Ag(s) I2(s) + 2 e- 2I-(aq) Cu2+(aq) + 2 e-Cu(s) RA OA Cl- Br2(l) OA RA Ag(s) I2(s) I-(aq) OA RA Cu2+(aq) RA OA SOA SRA

30. Remember… (Is Reduced) (Is Oxidized)

31. Predicting RedoxReactions (13.2b) • Now that you know what redox reactions are, you will be responsible for determining if a reaction will occur (is spontaneous) and if so, what the reaction equation will be. How do we do this? • The first step is to determine all the entities that are present. • Helpful reference: Table 6 pg. 575 • Remember: In solutions, molecules and ions behave independently of each other. • Example: When a solution of potassium permanganate is slowly poured through acidified iron(II) sulfate solution. • Does a redox reaction occur and what is the reaction equation?

32. Predicting Redox Reactions • The second step is to determine all possible OA’s and RA’s • This is a crucial step!! Things to watch out for: • Combinations • (i.e. MnO4-(aq) is an oxidizing agent only in an acidic solution) • To indicate this draw an arc between the permanganate and hydrogen ion • Species that can act as both OA and RA • Any lower charge multivalent metal i.e. Fe2+, Cu+, Sn2+, Cr2+ • Water (H2O(l)) • Label both possibilities in your list

33. Predicting Redox Reactions in Solution • We can use a redox table to predict which reactions will occur spontaneously. • If we assume that collisions are completely random, the strongest OA and the strongest RA will react. • Tips for using the redox table: • Choose the strongest OA and RA present in your mixture • Read reduction half-reaction equations from left  right • Read oxidation half-reaction equations from right  left • Assume any substances not present in the table to be spectator ions which do not need to be considered

34. Predicting Redox Reactions SOA • The third step is to identify the SOA and SRA using the data booklet • The fourth step is to show the ½ reactions (from the redox table) and balance: • SOA equation straight from table. SRA equation read from right to left • Are these equations balanced? Do the number of electrons lost = electrons gained • If not, multiply one or both equations by a number then add the balanced equations SRA

35. Predicting Redox Reactions 5. The last step is to predict the spontaneity. Does the net ionic equation represent a spontaneous or non-spontaneous redoxreaction?? If the SOA above  Spontaneous SRA?? If the SRA below  Nonspontaneous SOA

36. Predicting Redox Reactions #2 Could copper pipe be used to transport a hydrochloric acid solution? • List all entities • Identify all possible OA’s and RA’s • Identify the SOA and SRA • Show ½ reactions and balance • Predict spontaneity Since the reaction is nonspontaneous, it should be possible to use a copper pipe to carry hydrochloric acid

37. Disproportionation • The redox reactions we have covered so far have one reactant (OA) which removes electrons from a second reactant (RA) if a spontaneous reaction is to occur. Although the OA and RA are usually different entities, this is not a requirement. • A reaction is which a species is both oxidized and reduced is called disproportionation (aka autoxidation or self oxidation-reduction) • Occurs when a substance can act as either as oxidizing or reducing agent • Example: Will a spontaneous reaction occur as a result of an electron transfer from one iron(II) ion to another iron (II) ion? • No! Using the redox table and spontaneity rule, we see that iron(II) as an oxidizing agent is below iron(II) as a reducing agent, so the reaction is nonspontaneous

38. Disproportionation • Example #2: Will a spontaneous reaction occur as a result of an electron transfer from one copper(I) ion to another copper (I) ion? (see p. 828) Cu+(aq) + 1 e-Cu(s) Cu+(aq) Cu2+(aq) + 1 e- 2 Cu+(aq) Cu2+(aq) + Cu(s) • YES! Using the redox table and spontaneity rule, we see that copper(I) as an oxidizing agent is above copper(I) as a reducing agent. Therefore, an aqueous solution of copper(I) ions will spontaneously, but slowly, disproportionate into copper(II) ions and copper metal. (Note: See pg. 578 Ex.2 for another example.)

39. Summary: Five Step method for Predicting Redox Reactions • Step 1: List all entities present and classify each as a possible oxidizing agent, reducing agent, or both. Do not label spectator ions. • Step 2: Choose the strongest oxidizing agent as indicated in a redox table, and write the equation for its oxidation. • Step 3: Choose the strongest reducing agent as indicated in the table, and write the equation for its oxidation

40. Step 4: Balance the number of electrons lost and gained in the half-reaction equations by multiplying one or both equations by a number. Then add the two balanced half-reaction equations to obtain a net ionic equation. • Step 5: Using the spontaneity rule, predict whether the net ionic equation represents a spontaneous or nonspontaneousredox reaction.

41. Homework… • Textbook: • p. 575 #25 a-c • p. 579 #26 • p. 582 #4-7, 13 • Diploma Questions: • MC 2011 #12, 16, 18, 23 • MC 2009 #15, 19, 25

42. Redox Reactions using Half-Reactions • Remember: a redox reaction includes BOTH an oxidation (RA) and a reduction (OA). • So far we have predicted redox reactions when the ½ reaction was provided to us in the Redox table. But what if the table does not provide the half reaction? • We can use our own knowledge to create the equation Rules for Writing Half-Reactions • Write an unbalanced ½ reaction showing formulas for reactants and products • Balance all atoms except H and O • Balance O by adding H2O(l) • Balance H by adding H+(aq) • Balance the charge by adding e- and cancel anything that is the same on both sides For basic solutions only: • Add OH-(aq) to both sides to equal the number of H+(aq) present • Combine H+(aq) and OH-(aq) on the same side to form H2O(l). Cancel equal amounts of H2O(l) from both sides.

43. Practicing Half-Reactions • Copper metal can be oxidized in a solution to form copper(I) oxide. What is the half-reaction for this process? Cu(s) Cu2O(s) • Balance all atoms except H and O 2Cu(s) Cu2O(s) • Balance oxygen by adding water 2Cu(s) +H2O(l) Cu2O(s) • Balance hydrogen by adding H+(aq) 2Cu(s) +H2O(l) Cu2O(s) + 2H+(aq) • Balance charge by adding electrons 2Cu(s) +H2O(l) Cu2O(s) + 2H+(aq) + 2 e-

44. Practicing Half-Reactions • Chlorine is converted to perchlorate ions in an acidic solution. Write the half-reaction equation. Is this half-reaction an oxidation or reduction? Cl2(g) ClO4-(aq) • Balance all atoms except H and O Cl2(g) 2ClO4-(aq) • Balance oxygen by adding water Cl2(g) + 8H2O(l)  2ClO4-(aq) • Balance hydrogen by adding H+(aq) Cl2(g)+ 8H2O(l)  2ClO4-(aq) + 16H+(aq) • Balance charge by adding electrons Cl2(g)+ 8H2O(l)  2ClO4-(aq) + 16H+(aq) + 14 e- Cl2(g) + 8H2O(l)  2ClO4-(aq) + 16H+(aq) + 14 e- OXIDATION

45. Predicting Redox Reactions by Constructing Half-Reactions SUMMARY • Use the information provided to start two half-reaction equations. • Using the rules we just learned about half-reactions • Balance each half-reaction equation. • Multiply each half-reaction by simple whole numbers to balance electrons lost and gained. • Add the two half-reaction equations, cancelling the electrons and anything else that is exactly the same on both sides of the equation. For basic solutions only: • Add OH-(aq) to both sides to equal the number of H+(aq) present • Combine H+(aq) and OH-(aq) on the same side to form H2O(l). Cancel equal amounts of H2O(l) from both sides.

46. Predicting Redox Reactions by Constructing Half Reactions • Example: A person suspected of being intoxicated blows into this device and the alcohol in the person’s breath reacts with an acidic dichromate ion solution to produce acetic acid (ethanoic acid) and aqueous chromium(III) ions. Predict the balanced redox reaction equation. • Create a skeleton equation from the information provided: • Separate the entities into the start of two half-reaction equations • Now use the steps you learned for writing half reactions • Now, balance the electrons lost and gained, and add the half reactions. Cancel the electrons and anything else that is exactly the same on both sides of the equation. Note: a similar example is on p.581 (in basic solution)

47. Homework… • Textbook: • p. 582 #14-16 • Diploma: