1. SOLUTIONS
Homogeneous Mixtures

3. Solution - Definition
A solution is a homogeneous mixture of two or more substances (elements or compounds) in no definite ratio by mass.
All homogeneous mixtures are solutions.
Solutions can exist in any phase
Liquid solutions must be clear (light can pass through) and can be colored.

4. Solvents and Solutes
An aqueous solution contains dissolved substances.
2 components of a solution: solvent and solute.
Solvent dissolves the solute; solute is what is being dissolved.

5. Question Which statement describes KCl(aq)?
KCl is the solute in a homogeneous mixture
KCl is the solute in a heterogeneous mixture
KCl is the solvent in a homogeneous mixture
KCl is the solvent in a heterogeneous mixture

6. Solution process

7. Electrolytes and Nonelectrolytes
An electrolyte is a substance that conducts an electrical current when in an aqueous solution or a molten state.
All ionic compounds are electrolytes.
Strong electrolytes are good conductors of electricity.
Weak electrolytes are weak conductors of electricity.
Nonelectrolytes do not conduct electricity when in an aqueous solution or in a molten state.

9. Question
Explain why CH4 is a nonelectrolyte.

10. Terminology Solute - substance that is being dissolved (lesser amount)
Solvent - substance that is doing the dissolving (greater amount)
Aqueous solutions - water solutions (water is the solvent)
KCl(s) -> K+(aq) + Cl-(aq)
Note the (aq) to indicate a water
solution

11. Terminology (cont.)
Solubility - the amount of solute that will dissolve in a given amount of solvent at a certain temperature.
Dissociation - breaking apart into its ions
KCl(s)  K+(aq) + Cl-(aq)
Miscible - mixable (water & alcohol)
Immiscible - will not mix (oil & water)

12. Solubility Factors
The following have an effect on how much of something can be dissolved-
Nature of the solute and solvent
Temperature
Pressure (gases)

13. “Like Dissolves Like”
In chemistry the phrase is used to describe how solutes and solvents will interact with each other.
In Both cases the substances are soluble
water
CCl4 liquid
ethanol
I2 solid
Both polar
Both nonpolar

14. Nature of the Substances
Solute
Type
Non Polar Solvent
Polar
Solvent
Nonpolar
Soluble
Insoluble
Ionic

15. Temperature
For most solids solubility in water increases with and increase in temperature.
Ex: dissolving sugar in tea
For gases the opposite is true; solubility in a liquid decreases with an increase with temperature. Why?
Increasing temp allows the gases to escape the liquid.

16. Pressure
Pressure has little no affect on the solubility of solids or liquids
Pressure affects solubility of a gas in a liquid. As pressure increases, the solubility of a gas increases.
Ex: Opening up a soda can.

17. Solubility Factors
The following will effect how fast something dissolves-
Particle size
Stirring

18. Particle Size
The rate at which a solute dissolves depends on the particle size of the solute. The bigger the particle the longer it takes to dissolve it.

19. Stirring
Stirring increases the rate at which a solute dissolves because during stirring the fresh solvent is continuously in contact with the solute.

20. Question
Describe two methods by which you could remove gases that are dissolved in water.

21. TEMPERATURE Refer to Table G of your Reference Tables for Chemistry
The solubility of most solids increases with an increase in temperature
The solubility of gases decreases with an increase in temperature

22. Temperature & Solubility
These substances are gases
Notice that the solubility of the gases decreases with an increase in temperature
The solubility of most solids (those on the table) increases with an increase in temperature
The solubility is given in grams of solute per 100 grams of water

23. Question
Which compound’s solubility decreases most rapidly as the temperature changes from 10oC to 70oC?
NH4Cl
NH3
HCl
KCl

24. Solution Conditions Supersaturated Grams solute/100 g H2O Unsaturated
Temperature(°C)

25. Solution Types
Unsaturated - the solution can dissolve more solute in the solvent at the specified temperature.

26. Solution Types
Saturated - the solution is holding as much solute as it can hold at the given temperature ( the rate of solution equals the rate of dissolution).

27. Solution Types
Supersaturated - the solution is holding more solute than it normally can hold at the specified temperature. These solutions are unstable and will seek to reach saturation when disturbed and the excess solute will precipitate out.
Example: rock candy

29. Question
How many grams of NaNO3 would be needed to saturate 200 grams of H2O at 40oC?

30. Units of Concentration (Table T)
ways to express the amount of solute in solution (concentration)
Mass percent of solute
Parts per million (ppm)
Molarity

31. Mass Percent of Solute
On Table T you will see the following relationship –
Mass % solute = (mass solute/total solution mass) x (100)
Example – In a solution prepared by dissolving 24 g of NaCl in 152 g of water, what is the % by mass of NaCl in solution?
Solution –
Mass % NaCl = (24 g/176 g) x 100 = 14

32. ppm On Table T you will see the following relationships
ppm solute = (mass solute/total solution mass) x 106
There is a relationship between mass percent and ppm
ppm = mass percent x 104

33. ppm Example
Example – In the United States and Canada, drinking water cannot contain more than x 10-4 mg of mercury per gram of sample of solution. In parts per million what would that be?
Solution
Ppm Hg = (5 x 10-4 mg Hg/1 x 103 mg) x = 0.5
Remember there are 1000 mg = 1 gram)

34. Questions
An aqueous solution has gram of oxygen dissolved in grams of water. In the space in your answer booklet, calculate the dissolved oxygen concentration of this solution in parts per million. Your response must include both a correct numerical setup and the calculated result.

35. Molarity (M) Remember moles (n) = given mass/molar mass
Formula from Table G
M = mols of solute /Liters of solution
Example – if 40 g of NaOH is dissolved in water to prepare 500 mL of solution, what is the molarity?
M = 1 mol/0.5 L = 2M (or a 2 molar solution)
Molar mass of NaOH = 40 g/ mol

36. # grams used (convert to moles first)
*Moles = mass/ molar mass
*Mass = mole X molar mass
Molarity Problems
Fill-in the following blanks
Substance
Molar Mass (g/mol)
# grams used (convert to moles first)
L of solution
Molarity (M)
KOH
56.1
100
2.0
CaCl2
111
0.5
NH4NO3 80
160
1.0
Ca(OH)2
74.1
1.5

37. Question
How many moles of solute are contained in 200 ml of 1M solution?

38. Percent by Volume
% by volume (v/v) = 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑥 100 Example: If a 10 mL of propanone is diluted with water to a total solution volume 200 mL, what is the percent by volume of propanone in the solution?

39. Question
A bottle of the antiseptic hydrogen peroxide H2O2 is labeled 3.0% (v/v). How many mL H2O2 are in ,L bottle of this solution?

40. Molarity Dilution Problems
The number of mols of solute in a sample remains constant when we add water to dilute the sample, the concentration changes.
M1V1 = M2V2
If a 200 mL of a 3.0 M HCl solution is diluted to 400 mL what is the new molarity?

41. Question
How many milliliters of a solution of 4.00M KI are needed to prepare ml of 0.760M KI?

42. Molecules vs. Ionic Compounds in Aqueous solutions
Some molecular compounds dissolve but do not dissociate into ions.
C6H12O6(s) (glucose)  C6H12O6 (aq)
1 mole of sugar gives 1 mole of sugar when dissolved in water

43. Molecules vs. Ionic Compounds in Aqueous solutions
Many ionic compounds dissociate into independent ions when dissolved in water
NaCl (s)  Na+(aq) + Cl-(aq)
1 mole of NaCl solid gives 1 mole of sodium ions and 1 mole of chloride ions when dissolves (total of 2 moles of dissolved particles).
MgCl2 (s)  Mg+2(aq) + 2Cl-(aq)
Total of 3 moles of dissolved particles

44. Colligative Properties
Colligative properties refers to properties of a solution that depend on the concentration of particles.
Vapor pressure
Boiling point
Freezing point

45. Vapor Pressure Lowering
The presence of any solute (salt or sugar) lowers the vapor pressure of the solvent.
The more moles of dissolved particles, the lower the vapor pressure.
The greater the concentration of a solute the more it lowers the vapor pressure.

46. Boiling Point Elevation
The presence of a nonvolatile solute (salt or sugar) raises the boiling point of the solvent.
Nonvolatile: does not easily vaporize
The greater the concentration of the solute, the more it raises the boiling point.
The more moles of dissolved particles, the higher the boiling point.
Boiling point elevates by 0.52 oC for each mole of particles in a kg of water.

47. Boiling Point (cont.)
The elevation of the boiling point depends on the number of mols of particles present. For each mol of particles in a kg of water the boiling point of the aqueous solution is elevated by 0.52°C. That is called the boiling point elevation constant for water (0.52°C/m).
Molecular compounds such as C6H12O6 do not break up into ions, so one mol of C6H12O6 would raise the boiling point 0.52°C.
Ionic Compounds such as NaCl and CaCl2 break up into ions in aqueous solution. One mol of NaCl actually yields two mols of ions, one mol of CaCl2 yields three mols of ions. They would elevate the boiling point more than one mol of C6H12O6.

48. Freezing Point Depression
The presence of any solute (salt or sugar) lowers the freezing point of the solvent.
The more moles of dissolved particles, the lower the freezing point.
The freezing point of an aqueous solution is lowered 1.86°C for each mol of particles per kg of water.
The greater the concentration of a solute the more it lowers the freezing point.

49. Question
How are the boiling and freezing points of a sample of water affected when a slat is dissolve in the water?
Boiling point decreases and freezing point depression
Boiling point decreases and freezing point increases
Boiling point increases and freezing point decreases
Boiling point increases and freezing point increases

50. Question
Which solution containing 1 mole of solute dissolved in100 grams of water has the lowest freezing point?
KOH (aq)
C6H12O6 (aq)
C2H5OH (aq)
C12H12O11 (aq)

51. Summary The addition of solutes of water: Lowers vapor pressure
Raises boiling point
Lowers freezing point
The more particles dissolved the greater change it has on the colligative properties.

52. Other things to remember
Low vapor pressure=high boiling point
High vapor pressure = low boiling point
The stronger the intermolecular forces the lower the vapor pressure/higher the boiling point.