1. Chapter Nine Chemical Reactions in Aqueous Solutions

2. Section 9.1 General Properties of Aqueous Solutions

3. Review A solution is a homogeneous mixture
Gas example: air
Liquid liquid: salt water
Solid example: brass
Solute: substance being dissolved
Typically lesser in quantity
Solvent: substance doing the dissolving
Typically greater in quantity

4. Types of Solutes: Electrolytes vs. Nonelectrolytes
Electrolyte: substance that when dissolved in water conducts electricity
Sodium Chloride (or table salt)
Has ions in solution (dissociation)
Nonelectrolyte: substance that when dissolved in water does NOT conduct electricity
Sucrose (or sugar)
Does NOT have ions in solution, but molecules

5. Electrolytes vs. Nonelectrolytes

6. Strong vs. Weak Electrolytes
All water-soluble ionic compounds will dissociate completely
Therefore, they are strong electrolytes (i.e. substances that completely dissociate)
There are only 7 molecular compounds that are also considered strong electrolytes
HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4

7. Strong vs. Weak Electrolytes
Most molecular compounds are weak electrolytes OR nonelectrolytes
Weak electrolytes produce some ions upon dissolving but exist mostly of molecules that aren’t ionized
Acids are electrolytes (they produce H+ ions)
HCl(g)  H+(aq) + Cl-(aq)
Bases are electrolytes (they produce OH- ions)
NH3(g)  NH4+(aq) + OH-(aq)

8. Strong vs. Weak Electrolytes
For acids/bases that are WEAK, the reaction goes in both directions simultaneously
HC2H3O2(l)  H+(aq) + C2H3O2-(aq)
“” reaction occurs in both
directions
Dynamic Chemical Equilibrium
A + B2  AB2

9. Strong Electrolyte, Weak Electrolyte, or Nonelectrolyte???

10. Classify if Strong Electrolyte, Weak Electrolyte, or Nonelectrolyte
Sucrose (C12H22O11)
Fructose (C6H12O6)
Sodium Citrate (Na3C6H5O7)
Potassium Citrate (K3C6H5O7)
Ascorbic Acid (H2C6H6O6)

11. Section 9.2 Precipitation Reactions

12. Precipitation Reactions
Reaction where a “precipitate” forms

13. Precipitation Reactions

14. Solubility
Maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature

15. Solubility Rules for Ionic Compounds

16. Molecular, Ionic, & Net Ionic Equations
Pb(NO3)2(aq) + NaI(aq) 
Ionic Equation: Shows equation with ions dissociated
Net Ionic Equation: Shows only what’s involved in the reaction
Removes “Spectator Ions”

17. Group Quiz #1
For the following reaction, correctly predict the products to write the balanced molecular equation. Then write the ionic equation and the net ionic equation.
Aqueous solutions of Lead Acetate and Calcium Chloride

18. Section 9.3 Acid-Base Reactions

19. Acid-Base Models Arrhenius Model: Bronsted Model:
Acids produce H+ ions
Bases produce OH- ions
Bronsted Model:
Acids are H+ donors (or proton donors)
Bases are H+ acceptors (or proton acceptors)

20. More about Acids and Bases

21. Acid-Base Neutralization
Reaction between an acid and base
Produce water (most of the time) and a salt (ionic compound)

22. Section 9.4 Oxidation-Reduction Reactions

23. Oxidation-Reduction Reactions
A.K.A. “Redox” Reactions
Chemical Reaction where electrons are being transferred from one reactant to another.

24. Example Redox Reaction
Consider Zn(s) + CuCl2(aq)  ZnCl2(aq) + Cu(s)

25. Example Redox Reaction

26. Some definitions Oxidation is loss of electrons
Reduction is gain of electrons
“OIL RIG”
Oxidizing Agent: species that causes oxidation
Takes the electrons
Reducing Agent: species that causes reduction
Gives the electrons

27. Oxidation Numbers A.K.A. Oxidation State (or charge)
Help us determine what elements were oxidized and reduced
In order to determine an element’s oxidation number, you must follow the guidelines on the next two slides:

28. Guidelines with Oxidation Numbers

29. Guidelines with Oxidation Numbers

30. Determining Oxidation Numbers
What is the oxidation number of each atom in the following:
SO2
NaH
CO32-
H2SO4

31. What is Oxidized and What is Reduced?
2Fe + 6HBr  3H2 + 2FeBr3
N2 + 3H2  2NH3
2KClO3  2KCl + 3O2

32. Group Quiz #2
What is the oxidation number for chlorine in the compound HClO4?
What species is the reducing agent in the following equation?
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
Does the following equation represent a redox reaction? Why?
2Mg(s) + O2(g)  2MgO(s)

33. Section 9.5 Concentrations of Solutions

34. Concentration
Measure of amount of solute dissolved in a certain amount of solvent or solution
More solute:
Concentrated
Less solute:
Diluted

35. Molarity (One type of concentration)
Molarity = moles of solute/ L of solution
A.K.A. molar concentration
Represented by “M” ex: 1.5 M
If you have exactly 1 L of 1.5 M glucose, it contains 1.5 moles of glucose

36. Example
Suppose you wanted to make a M solution of KMnO4 using a 25o.00 mL volumetric flask. How would you do this?

37. Preparing molar solutions

38. Group Quiz #3
You need to make 500. mL of a M solution of Sodium Hydroxide (NaOH). What mass of NaOH do you need to use?
What is the molar concentration (M) of a solution prepared by dissolving g of Copper Chloride (CuCl2) in water to yield a 1.50 L solution?

39. Dilution Preparing less concentrated solutions
Typically done by adding water to concentrated solution
Dilution formula: McVc = MdVd
C = concentrated
D = diluted

40. Dilution Examples
What volume in mL of a 1.20 M HCl solution must be diluted in order to prepare 1.00 L of M HCl?
How much water was added?

41. Number of Ions in Solution
Recall: Soluble Ionic Compounds dissociate completely (all ionize)
If you have M of KMnO4, then there is M of K+ and M of MnO4- (1:1 ratio between ions)
[ ] are usually used to show concentration
[KMnO4] = M, [K+] = M, [MnO4-] = 0.500M

42. Number of Ions in Solution
If you have soluble ionic compounds with ratios other than 1:1 for ions, use subscripts to determine ion concentration
Ex: Na2SO4
[Na2SO4] = 0.35 M,
[Na+] = 0.70 M,
[SO42-] = 0.35 M
Suppose you had a 1.55 L solution of this ionic compound. How many moles of each ion do you have? How many individual ions do you have?

43. Section 9.6 Aqueous Reactions and Chemical Analysis

44. Gravimetric Analysis Analytical technique based on mass
Uses percent composition
Ex: A g sample of an ionic compound containing chloride ions and unknown metal cations is dissolved in water and treated with excess AgNO3. If g of AgCl precipitate, what is the percent by mass of Cl in the original compound?

45. Acid-Base Titrations Process where
Solution of known concentration (standard solution) is added gradually to
Another solution of unknown concentration till
The reaction is complete
Equivalence point: # of moles of H+ ions equals # of moles of OH- ions
End point: Color change in solution (visually indicates the equivalence point)

46. Acid-Base Titrations

47. Examples
What volume of a M NaOH solution is needed to neutralize 25.0 mL of a M H2SO4 solution?
If it takes mL of M HCl solution to neutralize mL of Ba(OH)2, what is the molarity of the base?

48. One more example
What is the molar mass of a diprotic acid if 30.5 mL of M NaOH is required to neutralize a g sample?

49. Group Quiz #4
How many milliliters of a 1.89 M H2SO4 solution are needed to neutralize 91.9 mL of a M KOH solution?
Explain the difference between an endpoint and an equivalence point.