1. Reactions in Aqueous Solutions
Chapter 4

2. General properties

3. Solution

4. Electrolyte

5. Hydration

6. HCl(l) H+(aq) + Cl−(aq)
CH3COOH(aq) H+(aq) + CH3COO−(aq)

7. Review of Concepts
The diagrams here show three compounds AB2 (a), AC2 (b), and AD2 (c) dissolved in water. Which is the strongest electrolyte and which is the weakest? (For simplicity, water molecules are not shown.)

8. Precipitate reactions

9. Double-displacement reaction

10. Solubility Rules

11. Equations Pb2+(aq) + 2I−(aq) PbI2(s)
Pb(NO3)2(aq) + 2 KI(aq) PbI2(s) + 2KNO3(aq)
Pb2+(aq) + 2NO3−(aq) + 2K+(aq) + 2I−(aq) PbI2(s) + 2K+(aq) + 2NO3(aq)
Pb2+(aq) + 2I−(aq) PbI2(s)

12. Example 4.2 page 125 in textbook
K3PO4(aq) + Ca(NO3)2(aq)
Example 4.2 page 125 in textbook

13. Review of Concepts
Which of the diagrams here acuratly describes the reaction between Ca(NO3)2(aq) and Na2CO3(aq)? For simplicity, only the Ca2+ (yellow) and CO32− (blue) ions are shown.

14. Acid-base reactions

15. General Properties Acid base Sour taste Color changes in plant dyes
React with metals to produce H2 gas
React with carbonates and bicarbonates to produce CO2 gas
Aqueous acid solutions conduct electricity
Taste bitter
Feel slippery
Color changes in plant dyes
Aqueous base solutions conduct electricity

16. Brønsted Acid and Bases
Proton donor
Monoprotic
Diprotic
Triprotic
Proton acceptor

17. Strong acids Strong basesHI
HBr
HClO4
HCl
H2SO4
HNO3
NaOH
KOH
LiOH
RbOH
CsOH                    
Ca(OH)2
Ba(OH)2
Sr(OH)2
Strong acids/bases are strong electrolytes and will completely dissociate in water.

18. Review of Concepts
Which of the following diagrams best represents a weak acid? Very weak acid? Strong acid? The proton exists in water as the hydronium ion. All acids are monoprotic. (For simplicity, water molecules are not shown.)

19. Acid-Base Neutralization
Reaction between an acid and a base
Generally aqueous solutions result in water and a salt
Ex: HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
*this is a strong acid and strong base so they completely dissociate and the net ionic equation is
H+(aq) + OH−(aq) H2O(l)
Ex: HCN(aq) + NaOH(aq) NaCN(aq) + H2O(l)
*this is a weak acid and strong base so the acid does not completely ionize in water. When writing the ionic and net ionic equations you cannot break the weak acid apart! The net ionic equation is
HCN(aq) + OH−(aq) CN−(aq) + H2O(l)

20. Gas formation
Certain salts react with acids to produce gaseous products
HNO3 breaks down into H2O(l) + NO2(g) + NO(g)
H2CO3 breaks down into H2O(l) + CO2(g)
H2SO3 breaks down into H2O(l) + SO2(g)
NH4OH breaks down into H2O(l) + NH3(g)
H2S(g)
CO2(g)
H2(g)
If you get one of these as a product in your molecular equation, they immediately breakdown as above
Gasses do not ionize

21. Double Replacement Rxns Review
Driving Force
How do you recognize it?
Precipitate
You must memorize the solubility rules. Any compound formed from two ions can be recognized as soluble (written as separate ions) or as a precipitate (written as a molecule).
Gas formed
You must memorize the combinations that decompose into gases (there are 4). You must also memorize the gases that form. For example, when you H2SO3 as a product, you must know it decomposes into H2O and SO2 gas.
Weak electrolyte
You must memorize the short list of strong acids and strong bases so you will recognize all the weak acids and bases that dissolve, but do not dissociate into ions. The weak base ammonia, NH3, is in this category. It exits in water as NH3(aq) and only slightly forms the ions NH4+ + OH−

22. Oxidation-reduction reactions

23. Half-reaction OIL RIG Oxidation reaction Reduction reaction
Reaction that involves the loss of electrons
Contains reducing agent-donates electrons
Involves the gain of electrons
Contains oxidizing agent-accepts electrons

24. Oxidation Number
Charge of the atom would have in a molecule if electrons were transferred completely
Rules
Uncombined elements = 0
Neutral compounds sum = 0
Ion = ion charge (polyatomic ions sum to charge)
Exceptions
Hydrogen +1 w/ nonmetals, −1 w/ metals
Oxygen −2 except w/ fluorine (+2), in peroxides (−1)
Fluorine ALWAYS −1

25. More common oxidation numbers are in red.

26. Example
2 Mg(s) + O2(g)  2 MgO(s)

27. Types of Redox Reactions
Combination 2 Al(s) + 3 Br2(l) AlBr3(s)
Decomposition 2 NaH(s) Na(s) + H2(g)
Combustion C2H8(g) + 5 O2(g) CO2(g) + 4H2O(l)
Displacement
Hydrogen Ca(s) + 2 H2O(l) Ca(OH)2(s) + H2(g)
Metal Zn(s) + 2 HCl(aq) ZnCl2(aq) + H2(g)
Halogen Cl2(g) +2 KBr(aq) KCl(aq) + Br2(l)
Disproportionation 2 H2O2(aq) H2O(l) + O2(g)

28. Activity Series
For Halogens:
F2 > Cl2 > Br2 > I2

29. Elements most likely to undergo disproportionation

30. Concentration n M = V Molarity = moles of solute liters of solution
Ex: 1M KCl solution
KCl(s) K+(aq) + Cl−(aq)
H2O
Ex: 1M Ba(NO3)2 solution
Ba(NO3)2(s) Ba2+(aq) + 2 NO3−(aq)
H2O

31. Example
How many grams of potassium dichromate (K2Cr2O7) are required to prepare a 250 mL solution whose concentration is 2.16M?

33. Dilutions
MiVi = MfVf

34. Example
Describe how you would prepare 5.00x102 mL of a 1.75M H2SO4 solution, starting with an 8.16M stock solution of H2SO4.

35. Review of Concepts
What is the final concentration of a 0.6M NaCl solution if its volume is doubled and the number of moles of solute is tripled?

36. Quantitative analysis
Gravimetric analysis
Titrations
Acid-base
redox

37. Gravimetric Analysis

38. Example
A g sample of an ionic compound containing chloride ions and an unknown metal is dissolved in water and treated with and excess of AgNO3. if g of AgCl precipitate forms, what is the percent by mass of Cl in the original compound?

39. Acid-base titrations

40. Example
How many mL of a 0.610M NaOH solution are needed to neutralize 20.0 mL of a 0.245M H2SO4 solution?

42. Redox titrations

43. Example
A mL volume of M KMnO4 solution is needed to oxidize mL of a FeSO4 solution in an acidic medium. What is the concentration of the FeSO4 solution in molarity? The net ionic equation is 5Fe2+ + MnO4− + 8H+  Mn2+ + 5Fe3+ + 4H2O