1. Chapter 13 Properties of Solutions

2. 13.1 Solutions
Solutions are homogeneous mixtures of two or more pure substances.
In a solution, the solute is dispersed uniformly throughout the solvent.
Those most solutions consist of a solid dissolved in a liquid, they also include liquids dissolved in other liquids or a gas dissolved in a liquid (see below).
Metals dissolved in another metal are alloys.

3. The Solution Process
The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

4. As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

5. If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

6. Energy Changes in Solution
The overall energetics of solution formation has three components:
Separation of solute particles
Separation of solvent particles
New interactions between solute and solvent
The overall enthalpy change to form a solution, ∆Hsoln, is the sum of three terms:
∆Hsoln = ∆H1 + ∆H2 + ∆H3

7. Solute-solvent interactions (∆H3) are exothermic. ∆H3 < 0
It takes energy to separate solute particles (∆H1) and energy to separate solvent particles (∆H2).
∆H1 > 0, ∆H2 > 0
Solute-solvent interactions (∆H3) are exothermic.
∆H3 < 0
If the attractive interactions between solute and solvent has a higher enthalpy change than the enthalpy changes for separating the solute and solvent particles, the overall reaction is exothermic. ∆Hsoln < 0
If the attractive interactions between solute and solvent has a lower enthalpy change than the enthalpy changes for separating the solute and solvent particles, the overall reaction is endothermic. ∆Hsoln > 0

8. To sum up:
The enthalpy change of the overall solution process (∆Hsoln ) depends on H for each of these steps.
The solution process can be either endothermic or exothermic.

9. A solution will not form if ∆Hsoln is too endothermic. Why?
Processes that are exothermic (i.e., the energy of the system decreases) tend to occur spontaneously.
A solution will not form if ∆Hsoln is too endothermic. Why?
The solvent-solute interaction must be strong enough to make ∆H3 comparable in magnitude to ∆H1 + ∆H2.
NaCl does not dissolve in gasoline (a nonpolar liquid) because the gasoline molecules would only experience weak dispersion forces with the NaCl ions. These interactions cannot compensate for the energies necessary to separate the ions.
Water (a polar liquid) does not dissolve in a nonpolar liquid (such as gasoline) because the strong hydrogen bonds between the water molecules cannot be overcome by any attractions of the water with the gasoline molecules.

10. However, some endothermic processes are spontaneous. Why?
Ex: NH4NO3 readily dissolves in water even though the overall process is endothermic.
There is an overall increase in randomness as the NH4+ and NO3- ions disperse throughout the solution.
Increasing the disorder or randomness (entropy) of a system, tends to lower the energy of the system.
The formation of solutions is favored by the increase in entropy (randomness) that comes from mixing.

11. Another example: the mixing of CCl4 and C6H14, both nonpolar compounds.
The increasing disorder or randomness (entropy) of the system, as the two liquids mix, tends to lower the energy of the system.
So, even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more disordered.

12. The solution process, therefore, involves two factors: a change in enthalpy and a change in entropy (randomness).
A solution will form unless solute-solute or solvent-solvent interactions are too strong compared to solute-solvent interactions.

13. Beware!
Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.
Dissolution is a physical change—you can get back the original solute by evaporating the solvent.
Chemical reactions, which produce different products form the reactants, sometimes look like dissolutions.

14. 13.2 Types of Solutions Saturated
Solvent holds as much solute as is possible at that temperature.
Dissolved solute is in dynamic equilibrium with solid solute particles.

15. Unsaturated
Less than the maximum amount of solute for that temperature is dissolved in the solvent.

16. Supersaturated
Solvent holds more solute than is normally possible at that temperature.
These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

17. 13.3 Factors Affecting Solubility
1. Solute-Solvent Interactions
Chemists use the axiom “like dissolves like”:
Polar substances tend to dissolve in polar solvents.
Nonpolar substances tend to dissolve in nonpolar solvents.

18. The more similar the intermolecular attractions, the more likely one substance is to be soluble in another. Ex: ethanol is completely miscible in water.

19. Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has dispersion forces) is not.

20. Vitamin A is soluble in nonpolar compounds (like fats).
Vitamin C is soluble in water.

21. Gases in Solution
In general, the solubility of gases in water increases with increasing mass.
Larger molecules have stronger dispersion forces.

22. 2. Pressure Effects
The solubility of liquids and solids does not change appreciably with pressure.
The solubility of a gas in a liquid is directly proportional to its pressure.

23. Henry’s Law Sg = kPg where Sg is the solubility of the gas (as M);
The relationship between pressure and the solubility of a gas.
Sg = kPg
where
Sg is the solubility of the gas (as M);
k is the Henry’s law constant for that gas in that solvent;
Pg is the partial pressure of the gas above the liquid.

24. 3. Temperature
Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

25. The opposite is true of gases:
Carbonated soft drinks are more “bubbly” if stored in the refrigerator.
Warm lakes have less O2 dissolved in them than cool lakes.