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… 1- 2- 3- 1+ Bromate Chlorate Iodate Carbonate Nitrate
Acetate
Azide
Cyanide
Hydroxide
Permanganate
Thiocyanate 2-
Carbonate
Chromate
Sulfate
Oxalate
Dichromate
Peroxide 3-
Arsenate
Borate
Phosphate
Citrate 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  NaNO 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