1. Solutions

2. Occur in all phases
The solvent does the dissolving, usually is the substance that is present in the greatest amount.
The solute is dissolved.
There are examples of all types of solutes dissolving all types of solvent (9 types).
We will focus on aqueous solutions.

3. Formation of a Solution
When 1 or more substances disperses uniformly throughout a solution.
We talked about intermolecular forces, they can also exist between solute and the solvent particles in a solution.
For example, If you were to dissolve solid KCl in water, it would readily dissolve because the attractive interactions between the ions and the

4. Polar water molecules are strong and overcome the lattice energy holding KCl together. The force of the ion-dipole attraction is strong enough to pull the ions out of the crystalline structure.
Solvation, when the particles of the solute are completely surrounded by the solvent. (called hydration when the solvent is water)

5. Forming a solution
To form a solution, not only do we need to break apart the bond (NaCl), but between water molecules also so there is room for the ions to fit in.
The enthalpy change is made of 3 energies.
ΔHsoln= ΔH1 + ΔH2 + ΔH3

6. Energy of Making Solutions
Heat of solution ( DHsoln ) is the energy change for making a solution.
Most easily understood if broken into steps.
1.Break apart solvent
2.Break apart solute
3. Mixing solvent and solute

7. 1. Break apart Solvent 2. Break apart Solute.
Have to overcome attractive forces. DH1 >0
2. Break apart Solute.
Have to overcome attractive forces. DH2 >0

8. 3. Mixing solvent and solute
DH3 depends on what you are mixing.
If the molecules are attracted to each other DH3 is large and negative (ion dipole).
If the molecules are not attracted to each other, then DH3 is small and negative (non polar solute- polar solvent).

9. Exothermic/ Endothermic
When calculating the enthalpy of solution, the sum can be + or -, depending on substances.
(-) exothermic, the solution would get warmer
(+) endothermic, the solution would get colder

10. Size of DH3 determines whether a solution will form
Solute and Solvent
DH3
Solution
DH2
DH3
Energy
Solvent
DH1
Reactants
Solution

11. Ice packs, heat packs (-) heat of solution, solution gets warm
(+) heat of solution, the solution gets cool

12. Spontaneity Exothermic processes are usually spontaneous
A solution will not form if it is too endothermic.
NaCl would not dissolve in gasoline, gasoline is non-polar, and the weak forces you would get between the non-polar gas and NaCl would not be enough energy to break apart the bond in NaCl

13. Like dissolves Like That’s where this saying came from.
In order to understand it, you must understand the relative strength of bonds.
When non-polar substances are mixed, they are all London Dispersion Forces, so there is no net energy change, however this is also a spontaneous process?

14. There must be something besides energy then that accounts for whether a reaction will proceed spontaneously or not. Two things: Enthalpy and Entropy

15. Types of Solvent and solutes
If DHsoln is small and positive, a solution will still form because of entropy.
There are many more ways for them to become mixed than there is for them to stay separate.

16. Spontaneity Two ways a reaction will proceed spontaneously:
1) Exothermic processes are usually spontaneous, the change tends to lower the energy of the system
2) Processes in which the disorder of the system increases also tend to occur spontaneously.

17. In sum
The formation of solutions is favored by an increase in disorder. A solution will form unless the solute-solute or solvent-solvent interactions are too strong compared to the solute-solvent interactions.

18. Physical / Chemical Solutions
Be careful to distinguish whether the formation of a solution is chemical or physical. If you can recover the original substance, than it is a physical change (dissolving in water).

19. Structure and Solubility
Water soluble molecules must have dipole moments -polar bonds.
To be soluble in non polar solvents the molecules must be non polar.

20. Saturated Solutions
When a solution is in contact with an undissolved solute, two opposing reactions are occurring, dissolving, and crystallization.
When the rate is equal, equilibrium is established, there is no further increase in the amount of solute that will dissolve.

21. The system is said to be saturated
Adding more solute to the system will not result in an increase in the concentration of the solution, the solute particles will remain undissolved
The amount of solute needed to form a saturated solution in a given quantity of solvent is known as the solubility of that solvent.

22. Unsaturated solutions have the capacity to hold more solute
Supersaturated solutions are sometimes possible to form. When dissolve the substance in hot solvent and cool it, all of the solute may remain dissolved even though the solubility has decreased with the decrease in temp.

23. Supersaturated solutions
Extremely unstable
Adding a seed crystal will result in the excess solute crystallizing out of the solution.

24. Factors affecting solubility
Natural tendency to move towards disorder
Attraction between solute and solvent, the stronger the attraction between solute and solvent, the greater the solubility of the solute in that substance.

25. Miscible/ immiscible
Liquids that mix in all proportions are said to be miscible
Liquids that do not mix are immiscible
(Hydrocarbons do not mix in water)

26. Solubility of alcohol in water
Alcohols have the OH end which is polar
Solubility of alcohols decrease with increasing mass because the chain is becoming more like a hydrocarbon
However if the number of OH groups increases, the solubility will increase
Network solids are not soluble in either polar or non polar solvents because of the strong bonding forces within the solid

27. Pressure effects
Changing the pressure doesn’t effect the amount of solid or liquid that dissolves
They are incompressible.
It does effect gases, the solubility of a gas in any solvent is increased as the pressure over the gas is increased.

28. Dissolving Gases
Pressure effects the amount of gas that can dissolve in a liquid.
The dissolved gas is at equilibrium with the gas above the liquid.

29. The gas is at equilibrium with the dissolved gas in this solution.
The equilibrium is dynamic.

30. If you increase the pressure the gas molecules dissolve faster.
The equilibrium is disturbed.

31. The system reaches a new equilibrium with more gas dissolved.
Henry’s Law.
P= kC
Pressure = constant x Concentration of gas

32. Try
Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of 2.0 atm over the liquid at 25°C. The Henry’s law constant for CO2 in water at this temperature is x mol/L-atm.

33. Temperature Effects
Increased temperature increases the rate at which a solid dissolves.
Usually it will increase the amount of solid that dissolves.

34. 20 40 60 80
100

35. Gases are predictable As temperature increases, solubility decreases.
Gas molecules can move fast enough to escape.

36. Concentration Dilute- small concentration of solute
Concentrate-large concentration of solid

37. 1. Quantitatively Mass percentage = mass solute x 100 mass solution
Very dilute in ppm = mass solute x 106
A solution of 1 ppm = 1g/ million g
Really dilute ppb= mass solute x 109

38. 2. Mole Fraction Moles component total moles of all components
Useful when dealing with gases

39. 3. Molarity
Molarity = moles of solute
liters of solution

40. 4. Molality Molality = moles of solute Kilograms of solvent
Molarity changes with temperature, molality does not.
Usually use molality when a solution is to be used over a range of temperatures.

41. Problems
Calculate the mass percentage of Na2SO4 in a solution containing 14.7 g of Na2SO4 dissolved in 345 g of water.
Calculate the mole fraction of methanol (CH3OH) when 7.5 grams of it is dissolved in 245 g of water.
Calculate the molarity of a solution when 10.5 g of KCL is dissolved in 250 ml of solution.
Calculate the molality of a solution when 13 g of benzene (C6H6) is dissolved in 17 g of CCl4.

42. Colligative properties
Add ethylene glycol to car radiators as antifreeze
Add salt to water to lower the freezing point
Also raises the boiling point.
Reduces the vapor pressure
Reduces osmotic pressure

43. Colligative Properties
Colligative properties depend only on the number - not the kind of solute particles present

44. Vapor Pressure
Vapor pressure is the pressure exerted by a gas above the liquid in a closed system at equilibrium
A non-volatile substance has no appreciable vapor pressure
A volatile substance would

45. Vapor Pressure of Solutions
A nonvolatile solute lowers the vapor pressure of the solution.

46. Psoln = csolvent x Psolvent
Raoult’s Law:
Psoln = csolvent x Psolvent
Vapor pressure of the solution = mole fraction of solvent x vapor pressure of the pure solvent
Doesn’t depend on the type of particle, only the concentration. Also, only works with non-volatile solutes otherwise they would be contributing to the pressure. (assume non-electrolytes)

47. Real Solutions Raoult’s Law- for ideal solutions
A solution only approaches ideal conditions when the solute concentration is low, and when the solute and solvent have similar molecular sizes and intermolecular forces.

48. Boiling point Elevation
Because a non-volatile solute lowers the vapor pressure it raises the boiling point.
The equation is: DT = Kbmsolute
DT is the change in the boiling point
Kb is a constant determined by the solvent.
msolute is the molality of the solute

49. Freezing point Depression
Because a non-volatile solute lowers the vapor pressure of the solution it lowers the freezing point.
The equation is: DT = Kfmsolute
DT is the change in the freezing point
Kf is a constant determined by the solvent
msolute is the molality of the solute

50. 1 atm
Vapor Pressure of pure water
Vapor Pressure of solution

51. 1 atm
Freezing and boiling points of water

52. 1 atm
Freezing and boiling points of solution

53. 1 atm
DTb
DTf

54. Electrolytes in solution
Since colligative properties only depend on the number of molecules, Ionic compounds should have a bigger effect.
When they dissolve they dissociate.
Individual Na and Cl ions fall apart.
1 mole of NaCl makes 2 moles of ions.
1mole Al(NO3)3 makes 4 moles ions.

55. Electrolytes have a bigger impact on omelting and freezing points per mole because they make more pieces.

56. Try
Automotive antifreeze consists of ethylene glycol C2H6O2, a non-volatile nonelectrolyte. Calculate the boiling point and freezing point of a 25. mass percent solution of ethylene glycol in water.
Kb= .52°C/m
Kf= 1.86°C/m

57. Try again!
List the following solutions in order of their expected freezing points: 0.05 m CaCl2, m NaCl, m HCl, HC2H3O2, m C12H22O11

58. Osmosis
Occurs through a semi permeable membrane. A membrane that usually allows the migration of small water molecules and not larger solute molecules.
The net movement of solvent is always toward the solution with a higher solute concentration.

59. Osmotic Pressure The pressure required to prevent osmosis (Fig 13.23).
Π= (n/V) RT=MRT
M=molarity
R= gas laws constant
T= temp (K)

60. Isotonic- when two solutions of identical osmotic pressure are separated by a semi-permeable membrane, no osmosis will occur
Hypotonic- when one solution has lower osmotic pressure compared to the other
The more concentrated one would be hypertonich

61. People who eat salty foods retain water
Cucumbers shrivel up in a salt water solution to make pickles
Osmosis is spontaneous.
Active transport-opposite

62. Try 2!
The average osmotic pressure of blood is 7.7 atm at 25°C. What concentration of glucose will be isotonic with blood?
What is the osmotic pressure at 20°C of a M sucrose C12H22O11 solution?

63. Colloids Somewhere between a solution and a heterogeneous mixture
Does not separate upon standing
Diameter of A
Scatter light- Tyndall Effect
When colloids are dispersed in water they can be hydrophyllic or hydrophobic