1. Chapters 1 – 3 Summary Measurements Basics of atomic structure
Chemical and molecular formulas
The concept of “moles”
Chemical equations and stoichiometry
Calculations involving stoichiometry

2. Chapter 4 Chemical Reactions
Ions in Aqueous Solutions
Types of Chemical Reactions
Working with Solutions
Quantitative Analysis

3. Ionic Theory of Solutions
Arrhenius Theory (1884):
Aqueous solutions that are electrically conducting contain ions.
(ie., The ions in aqueous solutions account for the electrical conductivity of the solution.)

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

6. A Solvent retains its phase (if different from the solute)
is present in greater amount (if the same phase as the solute)

7. Some Properties of Water
Water is “bent” or V-shaped.
The O-H bonds are covalent (e-‘s are shared).
Water is a polar molecule.
Hydration occurs when salts, strong acids and strong bases dissolve in water.

8. Overall charge on a molecule is “zero”.
Nuclear charge on hydrogen attracts negative charges
Electrons on oxygen attract positive charges
F = k (q1 * q2)/r12 2

9. Electrical conductivity of pure water
The light does not glow, thus no current goes through the solution.

10. A Solute dissolves in water (or other “solvent”)
changes phase (if different from the solvent)
is present in lesser amount (if the same phase as the solvent)

11. Ionic solid is solvated by water molecules

12. Solvated ions Solvated ions (ie., solutes) are called electrolytes.
Ionic solids and molecular compounds may both become solvated in water.
Electrolyte = a compound that dissolves in water to give an electrically conducting solution.

13. Electrolytes Strong - conduct current efficiently; highly ionized
NaCl, HNO3, HCl, NaOH
Weak - conduct a small current; medium ionization
vinegar (acetic acid), tap water
Non-electrolyte - no current flows; no ions formed in solution
pure water, sugar solution, CH3OH

14. Electrical Conductivity of a solution of NaCl
Thus, NaCl is highly ionized in water.

15. Electrical conductivity of ammonia
NH4OH is weakly ionized in water.

16. Reactions in Water NH3(aq) + H2O (l) NH4+(aq) + OH-(aq)kf
NH3(aq) + H2O (l) NH4+(aq) + OH-(aq)
Forward direction (strong electrolytes)
Reverse direction (weak electrolytes)
When kr is larger than kf, then the compound is weakly ionized. kr

17. Simple Rules for Solubility
1. Most nitrate (NO3-) 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.

19. Example
A precipitation reaction forming PbI2

20. 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)

21. 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.

22. Examples Molecular equation: Complete ionic:
Pb(NO3)2 + Na2SO PbSO4 + 2 NaNO3
Complete ionic:
Pb NO Na+ +SO42-
PbSO4 (s) + 2 Na+ + 2 NO3-
Net ionic:
Pb2+ (aq) + SO42-(aq) PbSO4 (s)

23. 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)
Oxidation-reduction reactions
Fe2O3(s) + Al(s) ® Fe(l) + Al2O3(s)

24. Acid-Base Concepts Two concepts of acid-base theory includes:
The Arrhenius concept: an acid increases the concentration of H+, and a base increases the concentration of OH-
Arrhenius acids/bases are only defined for water as a solvent.
The Brønsted-Lowry concept: an acid donates a proton, and a base accepts a proton.
The Bronsted-Lowry definition covers a wider range of compounds than the Arrhenius concept. 2

25. hydrochloric and sulfuric acid
Acids
Strong acids - dissociate completely to produce H+ in solution
hydrochloric and sulfuric acid
Weak acids - dissociate to a slight extent to give H+ in solution
acetic and formic acid

26. 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

28. Acid-Base Reactions Involving Gases
Molecular equation:
HNO3(aq) + KOH(aq) H2O(l) + KNO3(aq)
Total ionic: H+ + NO3- + K+ + OH H2O(l) + K+ + NO3-
Net ionic: H+(aq) + OH-(aq) H2O (l)

29. Performing Calculations for Acid-Base Reactions
1. List initial species and predict reaction.
2. Write balanced net ionic reaction.
3. Calculate moles of reactants.
4. Determine limiting reactant.
5. Calculate moles of required reactant/product.
6. Convert to grams or volume, as required.

30. Molarity
Molarity (M) = moles of solute per volume of solution in liters:
M = Molarity =
3 M solution of HCl = 3 moles HCl/liter soln
moles of solute
liter of solution

31. Concentration Example (Dilution)
How would one make 1 liter of a 3 M solution of HCl from a solution of concentrated HCl? (conc. HCl is 12 M)
use: M1 * V1 = M2* V2
or Mconc * Vconc = MHCl * VHCl
Vconc = = 0.25 liters = 250 ml
3M * 1 liter
12 M

32. Common Terms of Solution Concentration
Stock - routinely used solutions prepared in concentrated form.
Concentrated - relatively large ratio of solute to solvent. (5.0 M NaCl)
Dilute - relatively small ratio of solute to solvent. (0.01 M NaCl)

33. Volumetric Flask

34. Making solutions from solid solutes

35. Volumetric Flask

36. Add liquid up to the volume indicator mark
(ie., the meniscus)

37. Meniscus is
U-shaped in polar liquids,
Inverted U for non-polar liquids

38. Volumetric and Gravimetric Analysis (Titrations)
Methods used to accurately determine the concentration of a solution.
Gravimetric analysis involves precipitating a solute, and weighing the amount of precipitate formed (requires stoichiometric calculations).
Volumetric analysis involves using a solution of known concentration that reacts predictably and completely with the solution of unknown concentration.
(ie., How many ml of known solution are required to completely react with a solution of unknown concentration?)

39. Gravimetric Analysis

40. Acid-Base Indicators
Organic dyes that change color depending upon the amount of H+ (ie., H3O+) ions present in the solution.

41. Red Cabbage Extracts
Are Excellent
Acid-Base
Indicators

42. Volumetric Analysis: very similar the dilution calculations
It took 20 ml of a 3 M solution of NaOH to neutralize 50 ml of HCl.
What is the concentration of the HCl solution?
use: M1 * V1 = M2* V2
or MNaOH * VNaOH = MHCl * VHCl
MHCl = = 1.2 M
3 M * 20 ml
50 ml

43. Key Titration Terms
Titrant - solution of known concentration used in titration
Analyte - substance being analyzed
Equivalence point - enough titrant added to react exactly with the analyte
Endpoint - the indicator changes color so you can tell the equivalence point has been reached.

44. Titration Apparatus
Bruet
Titrant
Erlenmeyer Flask
Analyte

45. Final Volume of Titrant used
Indicator changes color at the endpoint

46. Titration Curves
Endpoint

47. pH curves and the buffer region
0 , because [base] = [acid]
buffer region
pH = pKa, at midpoint
of buffer region 2

48. Solution Examples NOTE:
Precipitation reactions are generally used for Gravimetric Analysis.
(ie., the sample is weighed to determine the amount of a substance in a sample).
A QUANTATIVE ANALYSIS METHOD
Problems:
76, 78, 80, 84, 105, 115

49. Oxidation-Reduction Reactions
A type of chemical reaction where an exchange of electrons occurs.
Oxidation – a loss of electrons (e-’s)
Reduction – a gain of electrons
Oxidizing Agent – receives electrons (is reduced)
Cu2+ (aq) + 2 e Cu (s)
Reducing Agent – donates electrons (is oxidized)
Fe (s) Fe2+ (aq) + 2 e-

52. Single Displacement Reaction
Cu (s)
2Fe (s) + 3CuSO Fe2(SO4)3 + 3Cu (s)
oxidized
reduced

53. Types of Redox Reactions
Combination Reactions
2 Na (s) + Cl2 (g) NaCl (s)
Decomposition Reactions
2 HgO (s) Hg (l) + O2 (g)
Single/Double Displacement Reactions
Zn (s) + 2 HCl (aq) ZnCl2 (aq) + H2 (g)
Combustion Reactions
2C4H10 (g) + 13O2 (g) CO2 + 10H2O (g)

54. Combination Reaction

55. Combustion Reaction

56. Decomposition Reaction

57. Rules for Assigning Oxidation States
1. Oxidation state of an atom in an element = 0
2. Oxidation state of monatomic element = charge
3. Oxygen = -2 in covalent compounds (except in peroxides where it = -1)
4. H = +1 in covalent compounds
5. Fluorine = -1 in compounds
6. Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion

59. Balancing by Half-Reaction Method (p. 625; sec. 19.1)
Write separate reduction, oxidation reactions;
2. For each half-reaction:
- Balance atoms other than H or O
Balance O using H2O
Balance H using H+
Make sure charges are balanced (use electrons to balance charges)
Use multipliers to balance e-’s

60. Balancing by Half-Reaction Method (continued)
3. If necessary, multiply by integer to equalize electron count.
4. Add half-reactions.
5. Check that elements and charges are balanced.

61. Half-Reaction Method - Balancing in Base
1. Balance as in acid.
2. Add OH- that equals H+ ions (both sides!)
3. Form water by combining H+, OH-.
4. Check elements and charges for balance.

62. Redox Reactions: Examples
Oxidation Numbers:
51, 53
Balancing Redox Equations: 60
Calculations involving Redox reactions:
134, 137