Understanding Chemical Reactions: Precipitation, Acid-Base, Redox

1. Chapter 4 Reactions in Aqueous Solutions

2. Types of Chemical Reactions • Chemical Reactions discussed in College Chemistry can be broken down into 3 main categories • Precipitation reactions • Acid-Base reactions • Oxidation-Reduction (redox) reactions

3. Types of Chemical Reactions • Precipitation Reactions:A process in which an insoluble solid (precipitate) drops out of the solution. • Clear solutions of two ionic compounds when mixed form a cloudy solution (cloudiness indicates solid)

4. Types of Reactions • Acid–Base Neutralization:A process in which an acid reacts with a base to yield water plus an ionic compound called a salt. • The driving force of this reaction is the formation of the stable water molecule.

5. Types of Reaction • Metathesis Reactions (Double Displacement Reaction): These are reactions where two reactants just exchange parts. AX + BY AY + BX

6. Types of Reactions • Oxidation–Reduction (Redox) Reaction:A process in which one or more electrons are transferred between reaction partners. • The driving force of this reaction is the decrease in electrical potential.

7. Precipitation Reactions • Develop the reaction equation • Balance the reaction equation • Predict the state of matter of each species present

8. Precipitation Reactions and Solubility Rules • To predict whether a precipitation reaction will occur must be able to predict whether a compound is soluble or not • Solubility rules

9. Solubility Rules • Salts - soluble: • All alkali metal and ammonium ion salts • All salts of the NO3–, ClO3–, ClO4–, C2H3O2–,and HCO3–ions

10. Solubility Rules • Salts which are soluble with exceptions: • Cl–, Br–, I– ion salts except with Ag+, Pb2+, & Hg22+ • SO42– ion salts except with Ag+, Pb2+, Hg22+, Ca2+, Sr2+, & Ba2+

11. Solubility Rules • Salts which are insoluble with exceptions: • O2– & OH– ion salts except with the alkali metal ions, and Ca2+, Sr2+, & Ba2+ ions • CO32–, PO43–, S2–, CrO42–, & SO32– ion salts except with the alkali metal ions and the ammonium ion

12. Precipitation Reactions and Solubility Rules • Predict the solubility of: • (a) CdCO3(b) MgO (c) Na2S (d) PbSO4(e) (NH4)3PO4(f) HgCl2

13. Precipitation Reaction • Precipitation reactions only occur if a solid is produced as a product. • If all products are aqueous compounds then no reaction has taken place.

14. Electrolytes in Aqueous Solution • To discuss aqueous reactions, must understand how different compounds interact with water • Electrolyte – substance that when added to water dissociates to form ions allowing for conductivity • Strong electrolytes – completely dissociate • Weak electrolytes – partially dissociate • Non-electrolyte – no dissociation in water

15. Aqueous Reactions and Net Ionic Equations • Molecular equations – substances involved in chemical reaction have been written using full formulas • Complete ionic equation – shows every species as it truly appears in the reaction vessel (aqueous strong electrolytes are shown as ions) • Net ionic equation – shows what is truly taking place in the reaction (all spectators are removed) • Must have a net ionic reaction for the reaction to occur as written

16. Complete Ionic Equation • AgNO3(aq) + KI(aq) → KNO3(aq) + AgI(s)

17. Net Ionic Equation • Shows species that undergo a change during the reaction • Obtained by removal of spectator ions.

18. Additional Examples • Reaction of magnesium acetate and sodium hydroxide in solution. • Reaction of lead nitrate with potassium dichromate in solution.

19. Acids, Bases and Neutralization Reactions • Acid / Base Definitions • Arrhenius • Acid – donates a H+ (H3O+) • Base – donates an OH- • Bronsted-Lowry • Acid – donates a H+ • Base – H+ acceptor

20. Acids, Bases and Neutralization Reactions

21. Electrolytes in Aqueous Solution • To discuss aqueous reactions, must understand how different compounds interact with water • Electrolyte – substance that when added to water dissociates to form ions allowing for conductivity • Strong electrolytes – completely dissociate • Weak electrolytes – partially dissociate • Non-electrolyte – no dissociation in water

22. Acids, Bases and Neutralization Reactions • Neutralization Reactions • Complete ionic reactions • Strong electrolytes (strong acids and bases) will dissociate entirely • Weak electrolytes are not shown as dissociated • Net ionic for strong acids reacting with strong bases will always be for the formation of water

23. Acids, Bases and Neutralization Reactions • Neutralization Reaction:produces salt & water. • HA(aq) + MOH(aq)  H2O(l) + MA(aq) • Write ionic and net ionic equations for the following: • (a) HBr(aq) + Ba(OH)2(aq) • (b) HCl(aq) + NH3(aq)

24. Neutralization Reactions • 2HBr (aq) + Ba(OH)2 (aq) → BaBr2(aq) + 2 H2O (l)

25. Neutralization Reactions • HCN (aq) + NaOH (aq) → NaCN (aq) + H2O (l)

26. Problem • Select the correct set of products for the following reaction.Ba(OH)2(aq) + HNO3(aq) →  • A.  BaN2(s) + H2O(l) • B.  Ba(NO3)2(aq) + H2O(l) • C.  Ba(s) + H2(g) + NO2(g) • D.  Ba2O(s) + NO2(g) + H2O(l) • E.  No reaction occurs

27. Problem • Select the net ionic equation for the reaction between lithium hydroxide and hydrobromic acid.LiOH(aq) + HBr(aq) → H2O(l) + LiBr(aq)  • a.  LiOH(aq) → Li+(aq) + OH-(aq) • b.  HBr(aq) → H+(aq) + Br-(aq) • c.  H+(aq) + OH-(aq) → H2O(l) • d.  Li+(aq) + Br-(aq) → LiBr(aq) • e.  Li+(aq) + OH-(aq)+ H+(aq) + Br-(aq) → H2O(l) + LiBr(aq)

28. Oxidation-Reduction Reactions • Redox reactions are those involving the oxidation and reduction of species (element or ion of an element). • Oxidation and reduction mustoccur together. They cannot exist alone. • Two important types • Single displacement reactions (activity series) • Combustions – reaction of a substance with O2

29. Oxidation Reduction Reactions Oxidation Is Loss (of electrons) AnodeOxidation Reducing Agent

30. Oxidation Reduction Reactions Reduction Is Gain (of electrons) Cathode Reduction Oxidizing Agent

31. Redox Reactions • Assigning Oxidation Numbers:All atoms have an “oxidation number” regardless of whether it carries an ionic charge. 1. An atom in its elemental state has an oxidation number of zero. 2. An atom in a monatomic ion has an oxidation number identical to its charge.

32. Redox Reactions 3. An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have if it were a monatomic ion. • A. Hydrogen can be either +1 or –1. • B. Oxygen usually has an oxidation number of –2. • In peroxides, oxygen is –1. • C. Halogens usually have an oxidation number of –1. • When bonded to oxygen, chlorine, bromine, and iodine have positive oxidation numbers.

33. Redox Reactions 4. The sum of the oxidation numbers must be zero for a neutral compound and must be equal to the net charge for a polyatomic ion. • A. H2SO4 2(+1) + (?) + 4(–2) = 0 net charge ? = 0 – 2(+1) – 4(–2) = +6 • B. ClO4– (?) + 4(–2) = –1 net charge ? = –1 – 4(–2) = +7

34. Problem • Sodium tripolyphosphate is used in detergents to make them effective in hard water. Calculate the oxidation number of phosphorus in Na5P3O10.  • A.  +3 • B.  +5 • C.  +10 • D.  +15 • E.  none of these is the correct oxidation number

35. Problem • The oxidation numbers of P, S and Cl in H2PO2-, H2S and KClO4 are, respectively  • A.  -1, -1, +3 • B.  +1, -2, +7 • C.  +1, +2, +7 • D.  -1, -2, +7 • E.  -1, -2, +3

36. Redox Reactions 5. Whenever one atom loses electrons (is oxidized), another atom must gain those electrons (be reduced). • A substance which loses electrons (oxidized) is called a reducing agent. Its oxidation number increases. • Asubstance which gains electrons (reduced) is called the oxidizing agent. Its oxidation number decreases.

37. Redox Reactions • For each of the following, identify which species is the reducing agent and which is the oxidizing agent. • Ca(s) + 2 H+(aq)  Ca2+(aq) + H2(g) • 2 Fe2+(aq) + Cl2(aq) 2 Fe3+(aq) + 2 Cl–(aq) • SnO2(s) + 2 C(s)  Sn(s) + 2 CO(g) • Sn2+(aq) + 2 Fe3+(aq)  Sn4+(aq) + 2 Fe2+(aq)

38. Problem • Identify the oxidizing agent in the following redox reaction.Hg2+(aq) + Cu(s) → Cu2+(aq) + Hg(l)  • A.  Hg2+(aq) • B.  Cu(s) • C.  Cu2+(aq) • D.  Hg(l) • E.  Hg2+(aq) and Cu2+(aq)

39. Problem • Sodium thiosulfate, Na2S2O3, is used as a "fixer" in black and white photography. Identify the reducing agent in the reaction of thiosulfate with iodine.2S2O32-(aq) + I2(aq) → S4O62-(aq) + 2I-(aq)  • a.  I2(aq) • b.  I-(aq) • c.  S2O32-(aq) • d.  S4O62-(aq) • e.  S2O32-(aq) and I-(aq)

40. The Activity Series of the Elements

41. The Activity Series of the Elements • Activity series looks at the relative reactivity of a free metal with an aqueous cation. • Fe(s) + Cu2+(aq)  Fe2+(aq) + Cu(s) • Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) • Cu(s) + 2 Ag+(aq)  2 Ag(s) + Cu2+(aq) • Mg(s) + 2 H+(aq) Mg2+(aq) + H2(g)

42. The Activity Series of the Elements Given the following three reactions, determine the activity series for Cu, Zn, & Fe. Fe(s) + Cu2+(aq)  Fe2+(aq) + Cu(s) Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) Fe(s) + Zn2+(aq)  NR

43. Optional Homework • Text – 4.24, 4.30, 4.32, 4.34, 4.36, 4.40, 4.42, 4.46, 4.48, 4.52, 4.54, 4.58, 4.60, 4.62, 4.64, 4.66, 4.68, 4.92, 4.100, 4.102, 4.108 • Chapter 4 Homework

44. Required Homework • Assignment 4