1. Chpt 11 - Solutions Concentrations Energy of solutions Solubility
Colligative Properties
Colloids
HW: Chpt 11 - pg , #s 12, 14, 20, 24, 29, 34, 41, 42, 44, 46, 52, 66, 72, 77, 81 Due Mon Dec. 3

2. Various Types of Solutions
Example
State of Solution
State of Solute
State of Solvent
Air, natural gas
Gas
Vodka, antifreeze
Liquid
Brass
Solid
Carbonated water (soda)
Seawater, sugar solution
Hydrogen in platinum
Solvent is majority component. Solute is minority component, usually the substance dissolved in the solvent (liquid).

3. Solution composition

4. Molarity
You have 1.00 mol of sugar in mL of solution. Calculate the concentration in units of molarity.
8.00 M
You have a 10.0 M sugar solution. What volume of this solution do you need to have 2.00 mol of sugar?
0.200 L

5. Molarity (M) example
Consider separate solutions of NaOH and KCl made by dissolving g of each solute in mL of solution. Calculate the concentration of each solution in units of molarity.
10.0 M NaOH
5.37 M KCl

6. Mass percent (%)
What is the percent-by-mass concentration of glucose in a solution made my dissolving 5.5 g of glucose in 78.2 g of water?
6.6%

7. Mole fraction (A)
A solution of phosphoric acid was made by dissolving 8.00 g of H3PO4 in mL of water. Calculate the mole fraction of H3PO4. (Assume water has a density of 1.00 g/mL.)
0.0145

8. Molality (m)
A solution of phosphoric acid was made by dissolving 8.00 g of H3PO4 in mL of water. Calculate the molality of the solution. (Assume water has a density of 1.00 g/mL.)
0.816 m

9. Solution Formation Schematic

10. Solution Formation Process
Separating the solute into its individual components (expanding the solute).
Overcoming intermolecular forces in the solvent to make room for the solute (expanding the solvent).
Allowing the solute and solvent to interact to form the solution.

11. Solution Formation Energies
Steps 1 and 2 require energy, since forces must be overcome to expand the solute and solvent.
Step 3 usually releases energy.
Steps 1 and 2 are endothermic, and step 3 is often exothermic.
Enthalpy change associated with the formation of the solution is the sum of the ΔH values for the steps:
ΔHsoln = ΔH1 + ΔH2 + ΔH3
ΔHsoln may have a positive sign (energy absorbed) or a negative sign (energy released).

12. Exo vs. Endo Hsoln
Demo NH4NO3 and NaOH examples

13. Explain why water and oil (a long chain hydrocarbon) do not mix
Explain why water and oil (a long chain hydrocarbon) do not mix. In your explanation, be sure to address how ΔH plays a role.
H1
H2
H3
Hsoln
Outcome
Polar solute, polar solvent
Large
Large, negative
Small
Solution forms
Nonpolar solute, polar solvent
Large, positive
No solution forms
Nonpolar solute, nonpolar solvent
Polar solute, nonpolar solvent

14. Solubility Factors Structural Effects: Pressure Effects:
Polarity (like dissolves like)
Pressure Effects:
Henry’s law (for dissolved gases)
Temperature Effects:
Affecting aqueous solutions

15. Pressure effects Henry’s law: C = kP
C = concentration of dissolved gas
k = constant
P = partial pressure of gas solute above the solution
Amount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution.

16. Gas solubility in liquid
Soda pop’s carbonated water has the carbon dioxide forced into the solution under pressure. When the can is opened Patm is much lower than Pcan so CO2 leaves -> pop goes flat.

17. Temperature effects
Although the solubility of most solids in water increases with temperature, the solubilities of some substances decrease with increasing temperature.
Predicting temperature dependence of solubility is very difficult.
Solubility of a gas in solvent typically decreases with increasing temperature.

18. Temp solubility charts

19. Colligative Properties
Depend only on the number, not on the identity, of the solute particles in an ideal solution:
Vapor pressure lowering
Boiling-point elevation
Freezing-point depression
Osmotic pressure

20. Vapor Pressure of solutions
If the Pvap of the solvent (water) > Pvap of the solution, equilibrium is reached when the solvent evaporates and the solvent is absorbed by solution. It does this to lower the Pvap towards its equilibrium value.

21. Raoult’s Law
Nonvolatile solute lowers the vapor pressure of a solvent.
Raoult’s Law:
Psoln = observed vapor pressure of solution
solv = mole fraction of solvent
Posolv = vapor pressure of pure solvent

22. Raoult’s Law - ideal solution
Ideal solution occurs with a nonvolatile solute in solution
Also the vapor pressure is then proportional to the mole fraction of the solvent using (total moles of ions of solute) in the solvent

23. Boiling Point elevation
Nonvolatile solute elevates the boiling point of the solvent.
ΔT = Kbmsolute
ΔT = boiling-point elevation
Kb = molal boiling-point elevation constant
msolute= molality of solute particles

24. Freezing Point depression
When a solute is dissolved in a solvent, the freezing point of the solution is lower than that of the pure solvent.
ΔT = Kfmsolute
ΔT = freezing-point depression
Kf = molal freezing-point depression constant
msolute= molality of solute particles

25. Phase Diagram of solutions

26. Boiling Point - Freezing Point explanation

27. Boiling Pt Elev Problem
A solution was prepared by dissolving g glucose in g water. The molar mass of glucose is g/mol. What is the boiling point of the resulting solution (in °C)? Glucose is a molecular solid that is present as individual molecules in solution. Kb = 0.51oC.kg/mol °C

28. Osmotic Pressure
Osmosis – flow of solvent into the solution through a semipermeable membrane. (Kidney dialysis uses this Principle).
= MRT
= osmotic pressure (atm)
M = molarity of the solution
R = gas law constant
T = temperature (Kelvin)

29. Osmotic Pressure graphic

30. Osmotic Pressure Problem
When 33.4 mg of a compound is dissolved in 10.0 mL of water at 25°C, the solution has an osmotic pressure of 558 torr. Calculate the molar mass of this compound.
Strategy: need Temp in K, Pressure in atm, use R with atm unit to get molarity. Then use vol get moles, then mass get molar mass.
111 g/mol

31. Colloids
Intermediate mixture - a heterogeneous mixture with particle size between a suspension and a solution
A suspension of tiny particles in some medium.
Tyndall effect – scattering of light by particles.
Suspended particles are single large molecules or aggregates of molecules or ions ranging in size from 1 to 1000 nm.

32. Types of colloids

33. Tyndall Effect graphic

34. Freezing Pt problem
You take 20.0 g of a sucrose (C12H22O11) and NaCl mixture and dissolve it in 1.0 L of water. The freezing point of this solution is found to be °C. Kf = 1.86oC.kg/mol
Assuming ideal behavior, calculate the mass percent composition of the original mixture, and the mole fraction of sucrose in the original mixture.
72.8% sucrose and 27.2% sodium chloride;
mole fraction of the sucrose is 0.313

35. Derivation of Colligative Properties
Specific derivation of the partial derivatives and derivation for colligative properties are found on the website.