Reaction Predictions

1. Key Points for Reaction Predictions

2. Review Nomenclature • Know symbols/names • Positive monatomic ions’ names are _____________. • Negative monatomic ions/names end in __________. • ____________ ion is named first.

3. Polyatomic Ions • Know “home base” ion (-ate ion) and charge. Others can be figured out from there. • “Home base” ion plus one oxygen ________________ • “Home base” ion minus one oxygen ________________ • “Home base” ion minus two oxygens ______________

4. You Gotta Know… • 3- • Arsenate • Borate • Phosphate • Citrate 1- • Bromate • Chlorate • Iodate • Nitrate • Acetate • Azide • Cyanide • Hydroxide • Permanganate • Thiocyanate 2- • Carbonate • Chromate • Sulfate • Oxalate • Dichromate • Peroxide • 1+ • Ammonium • Hydronium

5. Other Hints: • Bi- or hydrogen in the name means an extra hydrogen is present—Add H and add +1 to the charge. • Thio- means a sulfur replaces an oxygen—Add S & remove O; charge is the same.

6. Acid Names • Derive from negative monatomic or polyatomic ion • Monatomic ion plus hydrogen ________________ • -ate becomes ______ • -ite becomes ______

7. General Information • Uncombined elements have an oxidation number of zero. • Learn polyatomic ions and their charges!! • Think about reasonable possibilities for given reactants.

8. Equations • Word: Sodium chloride + silver nitrate  sodium nitrate + silver chloride • Formula: NaCl + AgNO3  NaNO3 + AgCl • Complete Ionic: Na + + Cl- + Ag+ + NO3-  AgCl + Na+ + NO3- • Net Ionic: Cl- + Ag+  AgCl

9. Net Ionic Equations • Dissociate all ions when in solution. (Is it soluble???) • Cancel ions that appear on both sides of an equation (spectator ions).

10. Solutions of nickel (II) nitrate and cesium hydroxide are mixed.

11. Equal volumes of equimolar sulfuric acid and sodium hydroxide are mixed.

12. Reaction Predictions on the a.p. exam

13. General Information • Will always appear on the exam • Three equations to write & three questions to answer • Must write a net ionic equation—NO SPECTATOR IONS • Balance the equations—atoms & charges. • All reactions will occur. Don’t worry about activity series.

14. Reaction Predictions One of a few general types: • Double Replacement or Metathesis • Single Replacement (also redox) • Redox • Synthesis/Decomposition • Complex Ions • Combustion

15. Double Replacement or Metathesis Starting point: two compounds, often in solution • Look for an acid/base reaction • Look for the formation of an insoluble compound (precipitate) • Both can happen simultaneously.

16. Acid/Base • A solution of hydrochloric acid is combined with a solution of sodium hydroxide. • A solution of hydrofluoric acid is combined with solid zinc hydroxide.

17. Precipitates • A solution of barium chloride is combined with a solution of potassium sulfate. • Solutions of cobalt II chloride and lithium sulfite are combined.

18. Both • Sulfuric acid solution in combined with solid calcium hydroxide.

19. Single Replacement Starting point: one compound (often in solution) and one uncombined element • Like will replace like in a reaction. • Oxidation numbers must change since an uncombined element’s oxidation # is always zero.

20. Single replacement • Fluorine gas is bubbled through a solution of potassium chloride. • A piece of solid aluminum is placed in a solution of copper II sulfate.

21. Solid lithium metal is added to water.

22. Aqueous solutions of oxalic acid (H2C2O4) and excess potassium hydroxide are mixed.

23. Synthesis/Decomposition • Synthesis—two reactants combine • Synthesis—starting point will be two separate substances—possibly elements • Decomposition—one reactant breaks apart • Starting point will be only one reactant—look for heat or electricity

24. Synthesis • Be familiar with diatomics • Sulfur and phosphorus occur as S8 and P4. • Nonmetal oxides + water make acids. • Metal oxides + water make bases.

25. Synthesis • Sulfur is burned in oxygen. • Sulfur (VI) oxide is added to water • Calcium oxide is added to water.

26. Decomposition • Know some special types: • Metal carbonates metal oxides & CO2 • Metal hydroxides  metal oxides and water • Metal chlorates metal chlorides and oxygen gas • Hydrogen peroxide  H2O & O2

27. Decomposition • Solid potassium chlorate is heated. • Solid aluminum oxide is heated. • An electric current is passed through molten sodium chloride.

28. Solid potassium chlorate is strongly heated.

29. Redox Several starting points: • Key compounds or ions: MnO4-, H2O2, Cr2O72-, HNO3 • Metals with multiple oxidation states: Sn2+, Sn4+, Cr2+, Cr3+, Cr6+, etc. • Acidic or basic conditions • Sometimes halogens, i.e. I-, IO-, IO2-, etc.

30. Redox • In redox, one reactant gains electrons and one reactant loses. • LEO says GER: • Losing electrons = oxidation • Gaining electrons = reduction • One process cannot occur without the other!

31. Reducing Agents • Cause reduction in something else while being oxidized themselves • Electron donors (any species that loses electrons) • Metal atoms • Negative ions • Positive metal ions that may still be able to lose more electrons

32. Oxidizing Agents • Cause oxidation in something else while being reduced themselves • Electron acceptors • Nonmetal atoms • Positive ions • Permanganates, peroxides and dichromates

33. Autooxidation • Some species such as peroxide are self-oxidzining • One species is both oxidized and reduced • Ex: H2O2  H2O + O2

34. Acidic or Basic Conditions • See page 833 • Reactions may be different under acidic or basic conditions • Use the reduction potential table for reference

35. To Predict: • Single replacement—follow normal pattern • Synthesis—follow normal pattern • All others: • Think about the most stable state of a substance (i.e. K vs K+) • Use the standard reduction potential table to help

36. An acidic solution of potassium dichromate is added to a solution of iron II nitrate • Find acidified dichromate as a reactant (oxidizing agent—It is reduced.) • Find iron II as a product. (It is oxidized.) • Reverse the one that is oxidized (iron II). • Combine and balance.

37. A strip of copper is immersed in dilute nitric acid • Copper begins with ox. # of zero; can only form positive ions; must be oxidized; reverse reaction. • NO3- + 4H+ NO + 2H2O

38. SUMMARY • Two uncombined reactants—synthesis • Single reactant—decomposition • Water as a reactant—metals & metal oxides produce bases; nonmetals and nonmetal oxides produce acids.

39. SUMMARY • Acid & base reactants (including salts)—neutralize to salt and water • Two salt solutions—look for a precipitate or a gas produced • Combustion of hydrocarbon—produce CO2 & H2O • Solid metal placed in solution—single replacement/redox rxn

40. SUMMARY • Transition metal with ammonia (NH3), hydroxide (OH-), cyanide (CN-), or thiocyanate (SCN-)—form complex ions in which ions attach to metals • Doesn’t matter how many—just get charge correct

41. Special Notes • If a product would be carbonic acid (H2CO3), it will break down into CO2 and H2O. • If a product would be ammonium hydroxide, it would break down into NH3 and H2O. • Ammonium carbonate decomposes into CO2, NH3 & H2O.

42. EXAMPLES • A piece of aluminum metal is added to a solution of silver nitrate. • Al + 3Ag+ Al3+ + 3Ag • A piece of solid bismuth is strongly heated in oxygen. • 4Bi + 3O2  2Bi2O3 (or another oxide)

43. EXAMPLES • An excess of sodium hydroxide solution is added to a solution of magnesium nitrate • Mg2+ + 2OH- Mg(OH)2 • Solid lithium hydride is added to water • LiH + H2O  LiOH + H2 • A concentrate solution of ammonia is added to a solution of zinc nitrate • Zn2+ + 4NH3  Zn(NH3)42+

44. Examples • Concentrate hydrochloric acid is added to solid manganese II sulfide. • 2 H+ + MnS  H2S + Mn2+ • Excess chlorine gas is passed over hot iron filings. • 3Cl2 + 2Fe  2FeCl3 (possibly FeCl2)

45. Examples • Solid ammonium carbonate is heated • (NH4)2CO3  CO2 + NH3 + H2O • Equal volumes of 0.1 M sulfuric acid and 0.1 M potassium hydroxide are mixed • H+ + OH-  H2O

46. EXAMPLES • Propanol is burned completely in air • 2CH3CH2CH2OH +9 O26 CO2 + 8H2O or 2C3H7OH +9 O26 CO2 + 8 H2O • A solid sample of magnesium carbonate is heated strongly. • MgCO3  MgO + CO2 • Ethene gas is bubbled through a solution of bromine. • C2H4 + Br2  C2H4Br2