1. Solutions

2. Aqueous Solutions
Water containing dissolved substances are aqueous solutions.
The dissolving medium is the solvent.
The dissolved particles are the solute.
Solutions are homogeneous mixtures.
Solute particles will pass through filter paper.

3. Dissolving of solid sodium chloride.
solute
solvent
solution

4. Polar water molecules interacting with positive and negative ions of a salt.

5. Solvation Solvent molecules attract solute particles.
Solute particles break away from each other and the crystal structure breaks apart.

6. Solvation
Miscible: two liquids can dissolve in each other. Ex. water and ethanol
Polar solvents will dissolve polar solutes.
Non-polar solvents dissolve non-polar solutes.
“Like dissolves like”

7. The ethanol molecule contains a polar O—H bond.

8. The polar water molecule interacts strongly with the polar O—H bond in ethanol.

9. An oil layer floating on water

10. Electrolytes
Compounds that conduct electric current in aqueous solutions or molten state.
All ionic compounds, including insoluble ones, are electrolytes.
Some molecular compounds can conduct electricity. (HCl, NH3)
Strong and weak electrolytes.

11. electrolyte
non-electrolyte

13. Properties of Solutions

14. Rate of Dissolution Agitation: increases rate of dissolution. (rate)
Temperature: increase in temperature increases dissolution. (rate and amount)
Particle size: decrease in particle size increases dissolution. (rate)

16. Solubility
Amount of a substance that dissolves in a given amount of solvent at a given temperature to produce a saturated solution Ex. sodium 0 oC  4.76 g/100 mL
Saturated solution: max. amount of solute that can be dissolved in a given amount of solvent at a specific temperature.

17. (a) A solid solute is added to a fixed quantity of water
(a) A solid solute is added to a fixed quantity of water. (b) After a few minutes, the solution is colored due to the dissolved solute, and there is less undissolved solute than in (a). (c) After a longer time, the solution color has deepened, and the quantity of undissolved solute is further diminished from that in (b). The solution in (b) must be unsaturated because more can dissolve. (d) Still later, the solution color and the quantity of undissolved solute appear to be the same as that in (c). Dynamic equilibrium must have been attained in (c) and persists in (d). In both (c) and (d), the solution is saturated.

18. Solubility
Supersaturated solution: a solution that contains more solute than it can theoretically hold at a given temperature.
If the solution is agitated or a seed crystal is placed in the solution crystallization will occur.

19. Supersaturated solution

20. Solubility of Liquids
Dynamic Equilibrium: number of particles dissolving is equal to the number of particles forming.

21. N2O4 and NO2 molecules are shown involved in bond breaking and bond forming

22. Concentration

23. Concentration/Molarity/Molar Concentration
A measure of the amount of solute that is dissolved in a measured amount of solvent.
Dilute: a solution that contains a small amount of solute .
Concentrated: a solution that contains a large amount of solute.

24. Concentration Units  mol/L or M
concentration = # of mol of solute (mol)
# of litres of solution (L)
C = n V
Units  mol/L or M

25. Examples c = n V c = 5.6 mol 20.8 L c = 0.27 M
What is the concentration of a solution if 5.6 mol of NaCl are dissolved to create a 20.8 L solution?
c = n V
c = 5.6 mol
20.8 L
c = M
WS 20/5 - 1

26. Parts per million
Very dilute concentrations can be recorded as parts per million. (ppm)
One ppm can be thought of as one drop in a bathtub of water.
1 ppm = 1 mg/L

27. Examples
What mass of chlorine is present in L of solution, if the solution is 6.00 ppm of chlorine?
1 ppm = 1 mg/L
6 ppm = 6 mg/L
6 mg x =
15.0 L
1 L
(6 mg) (15.0 L)
x =
= 90.0 mg
1 L

28. Examples
The label on a bottle of sparkling water lists the dissolved minerals as 440 ppm. What mass of minerals is present in a 200 ml glass of water?
1 ppm = 1 mg/L
440 ppm = 440 mg/L
440 mg x =
0.200 L
1 L
(440 mg) (0.200 L)
x =
= 88.0 mg
1 L

29. grams of solute dissolved in 100 mL solution
Percent solutions
Percent means parts per one hundred.
volume of solute
percent by volume =
X 100
volume of solution
mass of solute
percent by mass =
X 100
volume of solution
grams of solute dissolved in 100 mL solution

30. Examples
10.0 mL of acetic acid is diluted to a total volume of 200 mL. What is the percent by volume of acetic acid?
volume of solute
% (v/v) =
X 100
volume of solution
0.010 L =
X 100
0.200 L
= 5.00 % (v/v)

31. Preparing a solution

32. Convert 1.0 mol NaCl into grams of NaCl Obtain required mass of NaCl
Make a 1.0 M NaCl solution
Convert 1.0 mol NaCl into grams of NaCl
Obtain required mass of NaCl
Dissolve NaCl in 500 mL water
Transfer solution to volumetric flask
Add water to calibration line stopper and invert to mix

34. Dilution

35. Dilution
The process of making a concentrated solution into a less concentrated solution.
There are two types of dilution questions
simple dilution
addition dilution

36. Simple dilution: addition of water
n1 = n2
c1V1 = c2V2

37. c1V1 = c2V2

38. Dissociation Equations

39. Dissociation Equations
Non-electrolytes in solution separate into individual neutral, molecules that move freely throughout the solution. An equation showing a non-electrolyte dissolving simply shows it changing from the pure to its dissolved state.
C12H22O11(s) → C12H22O11(aq)

40. Dissociation Equations
When electrolytes dissolve they separate into ions that move freely throughout the solution. This is called dissociation. The solution as a whole remains neutral, since, although the individual ions are charged, they balance each other out.
Equations that show electrolytes dissolving show the solute changing from its pure state to aqueous ions.

41. Dissociation Equations
These equations are called dissociation equations.
Dissociation equations must be balanced, show correct ionic charge and physical states.
KCl(s) ↔ K+(aq) + Cl-(aq)
Al2(SO4)3(s) ↔ 2 Al3+(aq) + 3 SO42-(aq)
Cu(NO3)2(s) ↔ Cu2+(aq) + 2 NO3-(aq)

42. Calculating ionic Concentrations
In solution electrolytic compounds exist as free, separate ions.
NaCl(aq) really means Na+(aq) and Cl-(aq)
In chemical reactions involving such a solution the ions react independantly.
It is more correct to state the concentration of the ions present.

43. Calculating ionic Concentrations
To calculate the ionic concentration of ions in solution:
Step 1: write a balanced dissociation equation
Step2: use a mole ratio from the equation to determine the ion concentration

44. Calculating ionic Concentrations
What is the concentration of each ion in 0.23 M Al2(SO4)3 solution?
Balanced dissociation eqaution:
Al2(SO4)3(s) → 2 Al3+(aq) + 3 SO42-(aq)
Use mole ratio:
2 mol Al3+
0.23 M Al2(SO4)3
= M
1 mol Al2(SO4)3

45. Net Ionic Equations

46. Types of ionic Equations
Equations involving ionic compounds can be written three ways.
non-ionic equations
total ionic equations
net ionic equations

47. 2 AgNO3(aq) + BaCl2(aq) → 2 AgCl(s) + Ba(NO3)2(aq)
Non- ionic
The elements and compounds are written as molecules or formula units.
2 AgNO3(aq) + BaCl2(aq) → 2 AgCl(s) + Ba(NO3)2(aq)

48. Total ionic
Electrolytes are shown as separate dissociated ions while non-electrolytes, precipitates and gases are written as molecules or formula units.
2 Ag+(aq) + 2 NO-3(aq) + Ba2+(aq) + 2 Cl-(aq) → 2 AgCl(s) + Ba2+(aq) + 2 NO-3(aq)

49. Net ionic
Only the molecules, formula units or ions that have changed are shown in a net ionic equation.
Molecules and ions that do not change (spectator species) are not shown.
2 Ag+(aq) + 2 NO-3(aq) + Ba2+(aq) + 2 Cl-(aq) → 2AgCl(s) + Ba2+(aq) + 2 NO-3(aq)
2 Ag+(aq) + 2 Cl-(aq) → 2 AgCl(s)
Ag+(aq) + Cl-(aq) → AgCl(s)

50. Ion
Identification

51. Analysis for Metallic Elements
Qualitative analysis is designed to detect the presence of metal ions.
Quantitative analysis is designed to determine how much metal ion is present.

52. Solution Colour
The colour of a solution can be used to identify ions that are present in a solution.

53. What metals do colors indicate?
Solution colour
Ion present
colourless
Groups 1,2,17
blue
Cr2+
green
Cr3+
pink
Co2+
Cu+
Cu2+
pale green
Fe2+
Ni2+
purple
MnO4-

54. Flame test
The flame test is used to visually determine the identity of an unknown metal.

55. How is the test performed?
A clean platinum or nickel-chromium loop is required. They may be cleaned by dipping in hydrochloric acid, followed by rinsing with distilled water.
Test the cleanliness of the loop by inserting it into a Bunsen burner flame. If a burst of color is produced, the loop was not sufficiently clean.
The clean loop is dipped in a solution of an ionic salt. The loop with sample is placed in the clear or blue part of the flame and the resulting color is observed.

57. What metals do colors indicate?
Flame colour
Ion present
bright red
Li+
yellow
Na+
violet K+
yellow-red
Ca2+
bright-red
Sr2+
Ba2+
blue
Cu2+ (halide)
green
Cu2+ (others)
whitish-green
Zn2+

58. lithium
sodium
potassium

59. Selective Precipitation of Ions
Ions can be separated from each other based on their salt solubilities
Example: if HCl is added to a solution containing Ag+ and Cu2+, the silver precipitates while the Cu2+ remains in solution.
Removal of one metal ion from a solution is called selective precipitation.

61. Example
A precipitate is formed when HCl is added to a solution. Which of the following ions may be present?
silver ion
nickel ion
lead ion
calcium ion

62. Example
HCl, H2S, (NH4)3PO4 are added to a clear and colorless solution. No precipitate forms.Which of the following ions may be present?
cesium ion
nickel ion
lead ion
calcium ion