1. Aqueous solutions Types of reactions
Chapter 4
Aqueous solutions
Types of reactions

2. solution mixture homogeneous heterogeneous solvent Solute(s)
two phases
oil/water
milk
one phase
Tab water
solution
solvent
Solute(s)
more abundant
component of mixture
water in tab water
Less abundant or other component(s) of mixture
salts in tab water
Water always solvent even in 98% H2SO4

3. Solute undergoes Dissociation
solutes
Non-electrolytes
Electrolytes
Solute undergoes Dissociation
Sugar / H2O
strong
weak
complete
dissociation
partial dissociation
NaCl / H2O
HCl / H2O
HAc / H2O

4. Hydration of Solid Solute
At edges, fewer oppositely charged ions around
H2O can come in; Ion-dipole forces; Remove ion
New ion at surface
Process continues until all ions in solution
Hydration of ions
Completely surrounded by solvent

5. Molecular Compounds In Water
When molecules dissolve in water
Solute particles are surrounded by water
Molecules are not dissociated
Fig 5.6

6. Electrical conductivity of electrolyte solutions
Weak acids and bases
Ex. Acetic acid (HC2H3O2), ammonia (NH3)
Ionic compounds
Strong acids and bases
Ex. NaBr, KNO3, HClO4, HCl, KOH
Ex. Sugar, alcohol

7. Learning Check
Write the equations that illustrate the dissociation of the following salts:
Na3PO4(aq) →
Al2(SO4)3(aq) →
CaCl2(aq) →
Ca(MnO4)2(aq) →
3 Na+(aq) + PO43(aq)
Ca2+(aq) Cl(aq)

8. Solubility
Maximum amount of a substance that can be dissolved in a given amount of solvent at a given temperature.
Usually g/100 mL.
Saturated solution: Solution in which no more solute can be
dissolved at a given temperature
Unsaturated solution: Solution containing less solute than max.
amount; Able to dissolve more solute

9. Solubilities of Some Common Substances
Formula
Solubility
(g/100 g water)
Sodium chloride
NaCl
35.7 at 0°C
39.1 at 100°C
Sodium hydroxide
NaOH
42 at 0°C
347 at 100°C
Calcium carbonate
CaCO3
at 25°C

10. “Like dissolves Like” polar ↔ polar
Ethanol (C2H5OH) dissolves in water:
polar ↔ polar
Glucose (C6H12O6) and sucrose (C12H22O11) dissolve in water: polar ↔ polar
Oil doesn’t dissolve in water: nonpolar ↔ polar
Oil dissolves in benzene: nonpolar ↔ nonpolar

11. Water unable to separate Ag+ from Cl- Interaction very strong
Salts are polar.
soluble
insoluble
AgCl
NaCl
Water unable to separate Ag+ from Cl-
Interaction very strong

12. Relative Concentration
Solute-to-solvent ratio
Dilute solution
Small solute to solvent ratio
Ex. Eyedrops
Concentrated solution
Large solute to solvent ratio
Ex. Pickle brine
Dilute solution contains less solute per unit volume than more concentrated solution
Figure 5.2
Eyedrops = Low concentration of NaCl in water
Pickle brine = high concentration ofNaCl in water

13. Molarity quantitatively abbreviated M
1 M = 1 mol solute / 1 liter solution

14. Preparing Solution of Known Molarity
a b c d e
Weigh solid and transfer to volumetric flask
Add part of the water
Dissolve solute completely
Add water to reach etched line
Stopper flask and invert to mix thoroughly
Fig 5.24

15. Concentration of each type of ions in 0.50 M Co(NO3)2(aq)?
1 mol mol mol
In 1.00 L mol mol mol
Molarity M M M
Concentration of each type of ions in 0.50 M Fe(ClO4)3(aq)?
Molarity M M M

16. Moles of Cl- 1.75 L of 1.0×10-3 M ZnCl2(aq)?
1.75×10-3 mol ?

18. Practice
How many grams of HCl would be required to make mL of a 2.7 M solution?
What would the concentration be if you used 27g of CaCl2 to make 500. mL of solution? What is the concentration of each ion?
Describe how to make 1.00 L of a M K2CrO4 solution.
Describe how to make 250. mL of an 2.0 M copper (II) sulfate dihydrate solution.
Calculate the concentration of a solution made by dissolving 45.6 g of Fe2(SO4)3 to 475 mL. What is the concentration of each ion?

19. Describe how to make 1.00 L of a 0.200 M K2CrO4 solution.

20. No solid K2CrO4 available in the lab .
But 2.00 M K2CrO4 solution is available .

21. Dilution

22. Prepare 150 mL of 0.100 M H2SO4 from 16.0 M solution.
What volume of a 1.7 M solution is needed to make 250 mL of a 0.50 M solution?
18.5 mL of 2.3 M HCl is added to 250 mL of water. What is the concentration of the solution?
You have a 4.0 M stock solution. Describe how to make 1.0 L of a 0.75 M solution.

23. Types of Chemical Reactions
Reduction-Oxidation
Metathesis
Double Replacement
Electron transfer
AB + CD  AD + CB
precipitation reaction
Acid-Base Reaction
a solid is formed from solution
precipitate
Formation of a weak electrolyte
Formation of a gas

24. Precipitation reactions

25. Net ionic equation: describes what really happens.
Molecular equation
Ionic equation
Net ionic equation: describes what really happens.
Spectator ions:
A reaction takes place if it has a net ionic equation

26. Solubility Rules All nitrates and acetates are soluble
Salts of alkali metals ions and NH4+ ions are soluble.
Chlorides, bromides and iodides (salts of Cl-, Br- and I-) are soluble except those of Ag+, Pb2+, and Hg22+.
Most sulfates are soluble, except those of Pb2+, Ba2+, Hg2+, and Ca2+.
Most hydroxides are slightly soluble (insoluble) except those of alkali metals (Ba(OH)2, Sr(OH)2 and Ca(OH)2 are marginally soluble).
Sulfides (S2-), carbonates (CO32-), chromates (CrO42-) and phosphates (PO43-), are insoluble except those of alkali metals and NH4+.

27. Does the following mixing process involve a chemical reaction?

28. Precipitation reactions
NaOH(aq) + FeCl3(aq) ® ??
NaOH(aq) + FeCl3(aq) ® NaCl + Fe(OH)3
NaOH(aq) + FeCl3(aq) ® NaCl(aq) + Fe(OH)3(s)
Na+(aq)+OH-(aq) + Fe3+ (aq) + Cl-(aq) ® Na+ (aq) + Cl- (aq) + Fe(OH)3(s)
OH-(aq) + Fe3+ (aq) ® Fe(OH)3(s)

29. Precipitation reactions
BaCl2(aq) + KNO3(aq) ® ??
BaCl2(aq) + KNO3(aq) ® KCl + Ba(NO3)2
BaCl2(aq) + KNO3(aq) ® KCl(aq) + Ba(NO3)2(aq)
Ba2+(aq)+2 Cl-(aq) + K+ (aq) + NO3-(aq) ® K+ (aq) + Cl- (aq) + Ba2+(aq)+ 2 NO3-(aq)
No net ionic equation
No reaction

30. Practice iron (III) sulfate and potassium sulfide
Lead (II) nitrate and sulfuric acid.
solutions of NaOH and NiCl2 are mixed.

31.
0.15 mol ?

32. 1. 25 L of 0. 0500 M Pb(NO3)2 mixed with 2. 0 L of 0. 0250 M Na2SO4
1.25 L of M Pb(NO3)2 mixed with 2.0 L of M Na2SO4. Calculate the mass of precipitate.
mol ?
mol ?

33. Stoichiometry of Precipitation
What mass of solid is formed when mL of M Barium chloride is mixed with mL of M sodium hydroxide?
What volume of M HCl is needed to precipitate the silver from 50.0 ml of M silver nitrate solution ?
25 mL 0.67 M of H2SO4 is added to 35 mL of M CaCl2 . What mass CaSO4 Is formed?

34. HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2−(aq)
Arrhenius Acid
Substance that reacts with water to produce the hydronium ion, H3O+
Acid H2O  Anion + H3O+
HA H2O  A– + H3O+
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2−(aq) 
HCl(g) + H2O
Cl–(aq) + H3O+(aq)
Fig 5.10
Note for organic acids
In general, only hydrogen written first in formula transfers to H2O to give H3O+.
Bronsted-Lowry Acid: H+ donor

35. Bronsted Base: H+ acceptor
Arrhenius Bases
Substance that reacts with water to give OH–.
Metal Hydroxides
NaOH(s)  Na+(aq) + OH–(aq)
Mg(OH)2(s)  Mg2+(aq) + 2OH–(aq)
Basic Anhydrides
CaO(s) + H2O  Ca(OH)2(aq)
Ca(OH)2(aq)  Ca2+(aq) + 2OH–(aq)
c. Molecular bases:
NH3(aq)+H2O  NH4+(aq)+ OH-(aq)
Bronsted Base: H+ acceptor

36. Acid-Base Reactions Ionic equation
Weak electrolyte: H2O + H2O  H3O+(aq)+ OH-(aq)
Net ionic equation:
Any strong acid + strong base

37. HAc + H2O  H3O+(aq)+ Ac-(aq)
Formation of Weak electrolyte:
HAc + H2O  H3O+(aq)+ Ac-(aq)
Acid - Base Reactions are often called neutralization reaction Because the acid neutralizes the base.

38. Volume of 0.100 M HCl needed to neutralize 25.0 mL of 0.350 M NaOH ?
? ×10-3

39. 28.0 mL of 0.250 M HNO3 mixed with 53.0 mL of 0.320 M KOH;
Amount of water formed
Concentrations of H+ and OH- at the end of rct
7.0 mmol ? 
17.0 mmol ?

40. HNO3 is Limiting reactant: reacts completely No HNO3 left
HNO3 → H+ + NO3-
No H+ at the end of reaction
How much remains from KOH?
KOH → K+ + OH-
10 → mmol

41. Volumetric analysis: Titration
Controlled addition of 1 reactant to another until rxn is complete.
Acid-Base Titration: Very common type of titration
Ex. Analysis of citric acid in orange juice by neutralization with NaOH

42. Acid (Base) added equivalent to base (acid) present
An indicator is needed: organic substance that changes color according to solution acidity
Where the indicator changes color is the endpoint.
Endpoint must be very close to the equivalence point.
Acid (Base) added equivalent to base (acid) present
Phenolphthalein
Acidic Basic

43. Standardization of NaOH solution
Know the exact concentration!
Its weight is inaccurate .
NaOH is hygroscopic and it absorbs CO2.
Cannot be used to prepare solutions with exactly known M.
Not a primary standard.
KHP is a primary standard: high purity, no weighing problems,
Potassium hydrogen phthalate: KHC8H4O4.
Monoprotic!
41.2 mL of NaOH solution is needed to react exactly with g of KHP (M= g/mol).
MNaOH=?

46. practice
75 mL of 0.25M HCl is mixed with 225 mL of M Ba(OH)2 . What is the concentration of the excess H+ or OH- ?
A mL sample of aqueous Ca(OH)2 requires mL of M Nitric acid for neutralization. What is [Ca(OH)2 ]?