1. The Nature of Solutions

2. DEFINITIONS – SOLUTION – SOLUTE – SOLVENT – HOMOGENEOUS MIXTURE
CONCENTRATION – DILUTE VS CONCENTRATED
THE NATURE OF SOLUTIONS
SOLVATION

3. Some Definitions
A solution is a HOMOGENEOUS mixture of 2 or more substances in a single phase.
One constituent is usually regarded as the SOLVENT and the others as SOLUTES.

4. A. Definitions Solute - substance being dissolved
Solution - homogeneous mixture
Solute - substance being dissolved
Solvent - present in greater amount

5. Parts of a Solution
SOLUTE – the part of a solution that is being dissolved (usually the lesser amount). Uniformly spread in the solvent
SOLVENT – the part of a solution that dissolves the solute (usually the greater amount)
Solute + Solvent = Solution

6. What happens when a solute dissolves in a solvent?
Solvation – the process of dissolving
solute particles are surrounded by solvent particles
First...
solute particles are separated and pulled into solution
Then...

7. How Does a Solution Form?
As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

8. NaCl(s)  Na+(aq) + Cl–(aq)
B. Solvation
Dissociation
separation of an ionic solid into aqueous ions
Attractions between H2O and ions : molecule ion attractions
NaCl(s)  Na+(aq) + Cl–(aq)

9. CHARACTERISTICS OF A LIQUID SOLUTION
1.- Homogeneous mixtures, particles are evenly spread.
2.- Dissolved particles are too small to be seen, therefore solutions are clear and do not disperse light.
3.- Can not be separated by filtration. Dissolved particles are too small and will pass trough any filter.
4.- Stable. Dissolved particles will not come out of the solution and will not settle.

10. Solutions are not always liquids
Solutions are homogeneous mixtures of two or more pure substances.
In a solution, the solute (present in smaller amount) is dispersed uniformly throughout the solvent (present in largest amount).

11. Solutions Homogeneous mixtures.
Solvation is the process by which the solution forms.

12. SOLUBILITY FACTORS THAT AFFECT SOLUBILITY NATURE OF SUBSTANCES
TEMPERATURE
PRESSURE

13. Solubility
A measure of how much solute can be dissolved in an amount of solvent at a given temperature.

14. A substance can be…
Soluble in a solvent.
Example: sugar is soluble in water.
Miscible is the term used when the two components are liquids and they dissolve in one another.
Example: alcohol and water are miscible
Insoluble in a solvent
Example: sand is insoluble in water.
Immiscible is the term used when the two components are liquids and they do not mix.
Example: oil and water are immiscible

15. What affects Solubility?
1. Nature of Solute
2. Temperature
3. Pressure
* graph

16. “Like Dissolves Like” Nature of Solute
A polar solute molecule (alcohol) dissolves in a polar solvent (water).
A nonpolar solute (oil paint) dissolves in a nonpolar solvent (turpentine)
“Like Dissolves Like”

17. When temperature increases…
Solubility of a gas decreases
Solubility of a solid increases

18. 3 Pressure Makes gas more soluble
Has almost no effect on liquids and solids
Makes gas more soluble
 ex. Soda can
High pressure forces carbon dioxide into water to make soda.
- When you open the cap, there is less pressure on the soda b/c the soda fizzes and gas escapes.

19. Gases are more soluble at
high pressures
EX: nitrogen narcosis, the “bends,” soda

21. April 20
Objective: How do chemists represent solubility’s dependence on temperature?
Vocabulary:
Solubility
Saturated Solution
Unsaturated Solution

22. Do now Read and annotate FOCUS QUESTION
According to your reading, what is the difference between saturated and unsaturated solutions?

23. Solubility
A measure of how much solute can be dissolved in an 100g of solvent at a given temperature.

24. Review Concepts As temperature increases solubility
Increases for solid solutes
Decreases for gases solutes

25. Definitions Solutions can be classified as saturated or unsaturated.
A saturated solution contains the maximum quantity of solute that dissolves at that temperature. CAN NOT DISSOLVE MORE SOLUTE
An unsaturated solution contains less than the maximum amount of solute that can dissolve at a particular temperature.
CAN DISSOLVE MORE SOLUTE

26. Solubility Curves
Represent the solubility of different solutes in 100 g of water at different temperatures.
Each point on the curve represents a saturated solution at a given temperature.

27. C. Solubility Solubility Curve
shows the dependence of solubility on temperature
In the Y axis (dependent variable)
g solute/ 100g H2O
Remember
100g H2O = 100 ml H2O

29. Example 1 1 -How much NaNO3 can be dissolved at 10 o C in
100 g of water
The resulting solution is a saturated solution at 10 o C.

30. What happens if the temperature of the solution in increased from 10 oC to 50 oC?
Would that solution still be saturated?
What is the solubility of NaNO3 at 50 o C?
How much more solute can be dissolved?

31. Problem 6
Since the solubility of a solid increases as the temperature increases
Solubility at 50 o C ~115 g NaNO3/100 g H2O
Solubility at 10 o C ~ 80 g NaNO3/100 g H2O
115g-80g= 35 g
Additional 35 g of NaNO3 can be dissolved in the saturated solution at 10C to make it saturated at 50C

32. When the problem gives 2 temperatures get the solubility at the 2 different temperatures and analyze what would happen.

33. Example 2
What would happen if a 100 ml of a saturated solution of NH4Cl at 90 o C is cooled to 25 o C? (Hint find the solubility at each temperature!)

34. Example 3
How much NH4Cl is found in 200 ml of a saturated solution at 90 o C?
And in 50 mL at the same temperature?

35. Group Activity Get in your groups and work with the handout problems.
Each group is responsible to hand out one challenge problem (11 to 17) clearly showing the steps to solve it AND a sentence explaining the answer.
Homework
1 - Each student is responsible for 2 other problems of the challenge section (3 in total counting the one done in class).
2.-Read and annotate your RB P
Do questions 13 to 18 page 123

36. Exit Slip
How much KNO3 can be dissolved in 200g of H2O at 60 oC?

37. Solubility curves worksheet answers
1 KI
2 KClO3
3 SO2 g
5 SO2, NH3 and HCl
6 Sol at 50 o C ~115 g
at 10 o C ~ 80 g
difference 35 g
7 ~ 47 o C
8 KNO3 and NaNO3
9 NaCl
10 KNO3

38. Answer to challenge problems
11 Approximately 58 g of NaNO3 has to be added to 50 mL of water to make the solution saturated.
12 Approximately 44 g
o C
g will precipitate g g

39. Exit slip
Indicate what is the relationship between solubility and temperature for most of the substances in table G.
Why HCl, NH3 and SO2 do not follow the same relationship?
How much KNO3 can be dissolved in 100g of H2O at 60 oC?

40. Supersaturated solutions
Contain more solute that what is supposed to have a a given temperature.
They are VERY unstable!!!!

41. Supersaturated Sodium Acetate
One application of a supersaturated solution is the sodium acetate “heat pack.”

42. C. Solubility UNSATURATED SOLUTION more solute dissolves
no more solute dissolves
SUPERSATURATED SOLUTION
becomes unstable, crystals form
concentration

43. Solubility for ionic compounds Table F
This table is used to predict if a double replacement reaction will occur. If it the reaction produces an insoluble compound
it occurs. If the products of the reaction are filtered the insoluble compound will remain in the filter paper

44. Table F

45. Pb(NO3)2 + 2KI  PbI KNO3
NaCl + AgNO3  AgCl + NaNO3
CuSO4 + Na2CO 3  Na2SO4 + CuCO3

46. Questions
Is NaCl soluble?
Yes!
Is AgBr soluble?
No!

47. Set 1 Solubility curves 2 1 1(turn page around) 4(turn page again) 4 3
6 to 8 g
a- As P decreases, solubility decreases too.
b- As T increases, solubility decreases
12 a KNO3

48. April 23 Ways of expressing concentration MOLARITY
% BY MASS, BY VOLUME
PPM
Vocabulary
Concentration – dilute – concentrated
Molarity – volumetric flask
Homework: RB page q 24 to 37

49. Do Now: check answers to HW P 123 RB4 3 1
Page 120 1 2 3

50. The amount of solute in the solution.
CONCENTRATION
The amount of solute in the solution.
Relative terms
Diluted: Small amount of solute in relation to the amount of solvent
Concentrated: Large amount of solute in relation with the solvent.

51. Concentration of Solute
The amount of solute in a solution is given by its concentration.
Quantitative measurements are given by
Molarity ( M ) =
moles solute
liters of solution

52. Example 1 Molarity = moles of solute / liters of solution
If you have 50 moles (mol) of solute (salt) in 25 liters (L) of solution. Find the molarity (concentration).
Use Table T for the molarity formula.
Molarity = moles of solute / liters of solution
50 moles / 25 liters = 2M

53. Example 2
A solution contains 5 mol of NaCl in 1 Liter of solution. Find the molarity

54. Example 3
If a solution is 6 M , how many mol of solute are present in 1 L of solution?
How many mol in 2 L of solution?

55. Example 3 How to prepare 1L of a 0.1 M solution of CuSO4. 5H2O?
Step 1: Find the number of mol of the solute in the solution
Step 2. Convert mol to mass
Step 3. Place solute in volumetric flask
Step 4. Add water to the mark

56. Preparing Solutions
Weigh out a solid solute and dissolve in a given quantity of solvent.
Dilute a concentrated solution to give one that is less concentrated.

57. 1. 0 L of water was used to make 1. 0 L of solution
1.0 L of water was used to make 1.0 L of solution. Notice the water left over.

58. PROBLEM: Dissolve 5.00 g of NiCl2•6 H2O in enough water to make 250 mL of solution. Calculate the Molarity.
Step 1: Calculate moles of NiCl2•6H2O
Step 2: Calculate Molarity
[NiCl2•6 H2O ] = M

59. Calculating Concentrations
Dissolve 62.1 g (1.00 mol) of ethylene glycol in 250. g of H2O. Calculate molality and % by mass of ethylene glycol.

60. Learning Check
How many grams of NaOH are required to prepare 400. mL of 3.0 M NaOH solution?
1) 12 g
2) 48 g
3) 300 g

61. moles = M•V USING MOLARITY What mass of oxalic acid, H2C2O4, is
required to make 250. mL of a M
solution?
moles = M•V
Step 1: Change mL to L.
250 mL * 1L/1000mL = L
Step 2: Calculate.
Moles = ( mol/L) (0.250 L) = moles
Step 3: Convert moles to grams.
( mol)(90.00 g/mol) = g

62. Try this molality problem
25.0 g of NaCl is dissolved in mL of water. Find the molality (m) of the resulting solution.
m = mol solute / kg solvent
25 g NaCl mol NaCl
58.5 g NaCl
= mol NaCl
Since the density of water is 1 g/mL, 5000 mL = 5000 g, which is 5 kg
0.427 mol NaCl
5 kg water
= m salt water

63. APRIL 24 DO NOW – REVIEW MOLARITY % BY MASS, % BY VOLUME, ppm
COLLIGATIVE PROPERTIES
HW SOLUTIONS TAKE HOME TEST!!!!
DUE MONDAY

64. Two Other Concentration Units
% by mass
grams solute
grams solution
% by mass =

65. Two Other Concentration Units
Ppm = parts per million
grams solute
grams solution
Ppm =

66. Calculating Concentrations
Dissolve 62.1 g (1.00 mol) of ethylene glycol in 250. g of H2O. Calculate m & % of ethylene glycol (by mass).
Calculate weight %

67. Learning Check
A solution contains 15 g Na2CO3 and 235 g of H2O? What is the mass % of the solution?
1) 15% Na2CO3
2) 6.4% Na2CO3
3) 6.0% Na2CO3

68. Using mass %
How many grams of NaCl are needed to prepare 250 g of a 10.0% (by mass) NaCl solution?

69. Determining Electrical Conductivity
When a solution is soluble, it has ions that can conduct electricity (electrolytes)
Ex. NaCl
When a solution is insoluble, it cannot conduct electricity (non-electrolytes or poor electrolytes)
Ex. AgBr

70. Colligative Properties
On adding a solute to a solvent, the properties of the solvent are modified.
Melting point decreases
Boiling point increases
These changes are called COLLIGATIVE PROPERTIES.
They depend only on the NUMBER of solute particles relative to solvent particles, not on the KIND of solute particles.

71. Freezing point depression and boiling point elevation
One mole of particles dissolved in a 1000g of water lowers the freezing point of the water by 1.86 C and increases the boiling point of water by 0.52 C.
Note that electrolytes (ionic substances) produce a greater effect than non electrolytes because they produce more particles in solution.

72. Change in Freezing Point
Ethylene glycol/water
solution
Pure water
The freezing point of a solution is LOWER than that of the pure solvent

73. Change in Freezing Point
Common Applications of Freezing Point Depression
Ethylene glycol – deadly to small animals
Propylene glycol

74. Change in Freezing Point
Common Applications of Freezing Point Depression
Which would you use for the streets of New York to lower the freezing point of ice and why? Would the temperature make any difference in your decision?
sand, SiO2
Rock salt, NaCl
Ice Melt, CaCl2

75. Change in Boiling Point
Common Applications of Boiling Point Elevation

76. Freezing Point Depression
Calculate the Freezing Point of a solution containing 4.00 mol of glycol in a 1000 g of water
Kf = 1.86 oC/mol
Solution
Change in FP= (1.86 oC/mol)(4.00 mol)
∆TFP = 7.44
FP = 0 – 7.44 = oC (because water normally freezes at 0)
If NaCl is used instead the fp will be lower because the number of particles double.

77. Freezing Point Depression
At what temperature will a solution with 4 mol of NaCl in a 1000 g H2O freeze?
Solution
NaCl  Na + Cl
∆TFP = (1.86 oC/molal) • 8 mol
∆TFP = oC
FP = 0 – = oC

78. RB Answers Colligative Prop1 3 4 2 2 4 3 1