1. Reactions in Aqueous Solutions
Chapter 7

2. Water is the dissolving medium, or solvent.
Aqueous Solutions
Water is the dissolving medium, or solvent.

3. Driving Forces for Chemical Reactions
1. Formation of a precipitate (solid).
2. Formation of molecular substance (water).
3. Formation of a gas.
4. Transfer of electrons (Redox).

4. Types of Solution Reactions
Precipitation reactions
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Acid-base reactions
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)

5. pure H2O, sugar solution, glycerol
Electrolytes
Strong - conduct current efficiently - many ions in solution.
NaCl, KNO3, HNO3, NaOH
Weak - conduct only a small current - few ions in solution,
HC2H3O2, aq. NH3, tap H2O
Non - no current flows - no ions in solution.
pure H2O, sugar solution, glycerol

6. In what two ways can a solid ionic compound be made to conduct electricity?
Dissolve it in water.
Melt or fuse it.
Ions must be free to move
(mobile) in order to conduct electricity!

7. Dissociation ionic compounds
metal + nonmetal (Type I & II)
metal + polyatomic anion
Ammonium compounds
Acids
When ionic compounds dissolve in water the anions and cations are separated from each other; this is called dissociation
We know that ionic compounds dissociate when they dissolve in water because the solution conducts electricity 4

8. Dissociation
potassium chloride dissociates in water into potassium cations and chloride anions
KCl(aq) = K+ (aq) + Cl- (aq)
copper(II) sulfate dissociates in water into copper(II) cations and sulfate anions
CuSO4(aq) = Cu+2(aq) + SO42-(aq) K+
Cl- K Cl
Cu+2
SO42- Cu
SO4 5

9. K2SO4(aq) = 2 K+ (aq) + SO42-(aq)
Dissociation
potassium sulfate dissociates in water into potassium cations and sulfate anions
K2SO4(aq) = 2 K+ (aq) + SO42-(aq) K+
SO42- K
SO4 6

10. Ionic Compounds in Solution
In aqueous solution, soluble ionic compounds
exist in the form of ions.
K2CrO4(aq) + Ba(NO3)2(aq) ----> BaCrO4(s) + 2KNO3(aq)

11. Figure 7.1: The precipitation reaction that occurs when yellow potassium chormate, K2CrO4 (aq), is mixed with a colorless barium nitrate solution, Ba(NO3)2 (aq)

12. Solubility
1. A soluble solid readily dissolves in water-- designated with (aq).
2. A slightly soluble solid only dissolves to a tiny extent in water--designated with (s).
3. An insoluble solid does not dissolve to any appreciable extent in water--designated with (s).

13. Simple Rules for Solubility
1. Most nitrate (NO3), acetate (C2H3O2-), & chlorate (ClO3-) salts are soluble.
2. Most alkali (group 1A) salts and NH4+ are soluble.
3. Most Cl, Br, and I salts are soluble (NOT Ag+, Pb2+, Hg22+)
4. Most sulfate salts are soluble (NOT BaSO4, PbSO4, HgSO4, CaSO4)
5. Most OH salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble)
6. Most S2, CO32, CrO42, PO43 salts are only slightly soluble.

15. Precipitation Reactions
In all precipitation reactions, the ions of one substance are exchanged with the ions of another substance when their aqueous solutions are mixed
At least one of the products formed is insoluble in water
KI(aq) + AgNO3(aq)  KNO3(aq) + AgIs K+ I-
Ag+
NO3- Ag I 3

16. Describing Reactions in Solution
1. Molecular equation (reactants and products as compounds)
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
2. Complete ionic equation (all strong electrolytes shown as ions)
Ag+(aq) + NO3(aq) + Na+(aq) + Cl(aq) 
AgCl(s) + Na+(aq) + NO3(aq)

17. Describing Reactions in Solution (continued)
3. Net ionic equation (show only components that actually react)
Ag+(aq) + Cl(aq)  AgCl(s)
Na+ and NO3 are spectator ions.

18. Solutions of:
AgNO3
Na2SO4

19. If AgNO3 is mixed with Na2SO4 what ions are most abundant in the solution?
With what ions is the solution saturated?

20. 2AgNO3(aq) + Na2SO4(aq)  2 NaNO3 (aq) + Ag2SO4(s)
Molecular Equation
2Ag+(aq) + 2 NO3-(aq) + 2Na+(aq) + SO42-(aq)  2Na+(aq) + 2 NO3- (aq) + Ag2SO4(s)
Overall Ionic Equation
2Ag+(aq) + SO42-(aq)  Ag2SO4(s)
Net Ionic Equation
Silver Sulfate
Precipitate

21. Acids The nature of acids was discovered by Svante Arrhenius.
Acids are characterized by:
a sour taste (lemons -- citric acid).
producing H+ ions (protons) in aqueous solution.
turn blue litmus red.

22. Acids
Strong acids - dissociate completely (nearly 100 %) to produce H+ in solution
HCl, H2SO4, HNO3, HBr, HI, & HClO4
Weak acids - dissociate to a slight extent (approximately 1 %) to give H+ in solution
HC2H3O2, HCOOH, HNO2, & H2SO3

23. Figure 7.5: When gaseous HCl is dissolved in water, each molecule dissociates to produce H+ and Cl- ions

24. Bases Bases are characterized by: bitter taste (soap). feel slippery.
produce hydroxide (OH-) ions in aqueous solution.
turn red litmus blue.
A basic solution is said to be alkaline since they often contain one of the alkali metals --Na, K, Li, etc.

25. Bases Strong bases - react completely with water to give OH ions.
sodium hydroxide
Weak bases - react only slightly with water to give OH ions.
ammonia

26. Acid + Base ----> Salt + water
Acid-Base Reactions
Acid + Base ----> Salt + water
Molecular Equation
HCl(aq) + NaOH(aq) ----> NaCl(aq) + HOH(l)
Complete Ionic Equation
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) -----> Na+(aq) + Cl-(aq)
+ HOH(l)
Net Ionic Equation
H+(aq) + OH-(aq) -----> HOH(l)

27. Acid-Base Neutralization
The net ionic equation for the neutralization of any strong acid and base will always be:
H+(aq) + OH-(aq) -----> HOH(l)

28. Salts Salts are ionic compounds consisting of:
a. metal & nonmetal -- KCl
b. metal & polyatomic ion -- CuSO4
c. polyatomic ion & nonmetal -- NH4Cl
d. two polyatomic ions -- (NH4)2SO4

29. Oxidation-Reduction Reactions (Redox Reactions)
Redox reaction -- involves the transfer of electrons
loss and gain of electrons must be exactly equal.
loss and gain of electrons must be simultaneous.

30. Oxidation loss of electrons metal atoms -- Na, Ca, & K
Nao ---> Na1+ + 1e-
Cao ---> Ca2+ + 2e-
nonmetal ions -- Cl-, S2-, & O2-
Cl1- ---> Clo + 1e-
S2- ---> So + 2 e-

31. Reduction gain of electrons. nonmetal atoms. Oo + 2e- ---> O2-
Fo + 1e- ---> F1-
metal ions.
K1+ + 1e- ---> Ko
Ba2+ + 2e- ---> Bao

32. Figure 7.7: When powdered aluminum and iodine (shown in the foreground) are mixed (and a little water added), they react vigorously

33. Oxidation & Reduction Half-Reactions
Always add electrons (negative) to the more positive side of the equation.
The charge on both sides of an equation must be equal.
Oxidation -- Nao ----> Na e-
Reduction -- Cl2 + 2 e- ----> 2 Cl-

34. OIL RIG
Oxidation Is Loss.
Reduction Is Gain.

35. Redox Reactions Metal-nonmetal reactions are always redox reactions.
Any reaction that has a free element (such as O2) as a reactant or product are redox.
All single replacement reactions are redox.
All combustion reactions are redox.

36. Synthesis I (Composition)
A + X ----> AX ALWAYS REDOX
Element + Element -----> Binary Compound
Fe(s) + S(s) ----> FeS(s)
4 Al(s) + 3 O2(g) ----> 2 Al2O3(s)

37. Synthesis II (Composition)
A + X ----> AX NOT REDOX
Compound + Compound ----> Compound ( or more elements)
Ammonia + Acid ----> Ammonium Salt
NH3(g) + HCl(g) ----> NH4Cl(s)
2 NH3(aq) + H2SO4(aq) ----> (NH4)2SO4(aq)

38. Synthesis II (Composition) Continued
Water + An Oxide
Rule # 1 --Water + Metal Oxide ----> Metal Hydroxide (Base)
HOH(l) + CaO(s) ----> Ca(OH)2(s)
HOH(l) + Na2O(s) ----> 2 NaOH(aq)
Rule # 2 --Water + Nonmetal Oxide ----> Acid
HOH(l) + SO3(g) ----> H2SO4(aq)
HOH(l) + N2O5(g) ----> 2 HNO3(aq)

39. Single Replacement I A + BX ----> AX + B ALWAYS REDOX
Element + Compound ----> Different Element Different Compound
Metal Reactivities:
K > Na > Ca > Mg > Al > Zn > Cr > Fe > Cd > Co > Ni > Sn > Pb > H > Sb > Cu > Hg > Ag > Pt > Au

40. Single Replacement I (Continued)
A is a metal.
A is more reactive than B
A + BX ----> AX + B
Cu(s) + 2 AgNO3(aq) ---->Cu(NO3)2(aq) + 2 Ag(s)
Zn(s) + 2 HCl(aq) ----> ZnCl2(aq) + H2(g)
Rule # 3 -- Active Metal + Water ----> Metal Hydroxide + Hydrogen
Ca(s) + 2 HOH(l) ----> Ca(OH)2(s) + H2(g)

41. Single Replacement I (Continued)
A is a metal.
A is less reactive than B
A + BX ----> No Reaction (NR)
Cu(s) + HCl(aq) ----> NR

42. Single Replacement II Y + BX ----> BY + X ALWAYS REDOX
Element + Compound ----> Different Element Different Compound
Nonmetal Reactivities:
F > O > Cl > Br > I

43. Single Replacement II (Continued)
Y is a nonmetal.
Y is more reactive than X
Y + BX ----> BY + X
Cl2(g) + 2 KBr(aq) ----> 2 KCl(aq) + Br2(aq)
Y is less reactive than X
Y + BX ----> No Reaction
Cl2(g) + KF(aq) ----> NR

44. Single Replacement (Continued)
Remember!!
Metals replace metals.
Mg(s) + 2 Ag(NO3)2(aq) ----> 2 Ag(s) + Mg(NO3)2(aq)
Nonmetals replace nonmetals.
Cl2(g) + 2 KBr(aq) ----> Br2(aq) + 2 KCl(aq)

45. Figure 7.6: The thermite reaction gives off so much heat that the iron formed is molten

46. Decomposition I AX ----> A + X ALWAYS REDOX
AX is a binary compound.
Binary Compound ----> Element + Element
elect
2 NaCl(l) ----> 2 Na(l) + Cl2(g)
2 HOH(l) ----> 2 H2(g) + O2(g)

47. Decomposition II AX -----> A + X MAY OR MAY NOT BE REDOX
AX is a ternary compound.
Rule # 4 -- Metal Chlorates ----> Metal Chlorides Oxygen 
2 KClO3(s) ----> 2 KCl (s) + 3 O2(g) REDOX

48. Decomposition II (Continued)
Rule # 5 -- Metal Carbonates ----> Metal Oxides Carbon Dioxide 
CuCO3(s) ----> CuO (s) + CO2(g) NOT REDOX

49. Decomposition II (Continued)
Rule # 6 -- Metal Hydroxides ----> Metal Oxides Water 
Ca(OH)2(s) ----> CaO (s) + HOH(g) NOT REDOX

50. Decomposition II (Continued)
Rule # 7 -- Acid ----> Nonmetal Oxide + Water 
H2SO4(aq) ----> SO3(g) + HOH(g) NOT REDOX

51. Double Displacement (Double Replacement)
AX + BY ----> AY + BX NEVER REDOX
Compound + Compound ----> 2 New Compounds
Acid-Base & Precipitation Reactions (Chapter 7)
H2SO4(aq) + 2 NaOH(aq) ----> Na2SO4(aq) + 2 HOH(l)
NaCl(aq) + AgNO3(aq) ----> NaNO3(aq) + AgCl(s)

52. Combustion AX + O2 ALWAYS REDOX AX is a hydrocarbon. 
CxHy + O2(g) ----> CO2(g) + HOH(g)
CH4(g) + 2 O2(g) ----> CO2(g) + 2 HOH(g)
2 C6H6(l) + 15 O2(g) ----> 12 CO2(g) + 6 HOH(g)