1. Solute-Solvent Interactions
Ch. 13/14: Solutions
Solute-Solvent Interactions

2. “like dissolves like”
Rule of thumb for determining whether two substances will dissolve one another
Polar substances usually dissolve only polar
Nonpolar usually dissolve only nonpolar
Also includes type of bonding and types of intermolecular forces
Would ionic compounds be considered polar or nonpolar?

3. Hydration Charged ends of the water molecules attract ions in an
ionic compound
They surround the
ions and keep them away from other ions
As ions are pulled away from a crystal, more ions beneath are exposed
Eventually, all the ions are evenly dispersed throughout the solution

4. Hydrates
CoCl2.6H2O (cobalt chloride hexahydrate)
Ionic compounds that incorporate water molecules into their structure when they crystallize
Each have a specific number of molecules attached
What would happen if we crushed a hydrate and then heated it up?
CoCl2 (cobalt chloride, anhydrous)

5. Nonpolar Solvents Molecular compounds Ex:
carbon tetrachloride, CCl4
Toluene, C6H5CH3
Do not usually dissolve ionic compounds because they do not have a strong enough attraction to the ions to pull them apart

6. Miscibility Miscible Liquids that mix well in any proportion
Ex: benzene and carbon tetrachloride because they are both nonpolar and don’t have strong IMFs
Immiscible
Two liquids that are not soluble in each other
Ex: fats and oils in nonpolar liquids like toluene

7. Miscibility
Some substances have parts that are polar and nonpolar so are miscible in both types of liquids

8. Pressure and Solubility
Little effect on solubility of liquids and solids
Increase in P increases solubility of gases in liquids
Gas + solvent ↔ solution
When P is increased, it puts more gas molecules in contact with the surface of the liquid so more will be dissolved into the solution
Using Le Châtelier’s principle, as P is increased the reaction shifts towards the side of the reaction with less gas molecules

9. Pressure and Solubility

10. Henry’s Law
The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of that liquid at a constant T
Named after William Henry, English chemist
Effervescence
Quick escape of bubbles of gas from a liquid in which it was dissolved
In bottling plants, CO2 gas is forced into the soda with 5-10 atm of pressure

11. Gas Pressure and Solubility

12. Temperature and Solubility
Gases:
Increasing T usually
decreases gas solubility
As the molecules get more
KE, more can escape into
gas phase
Solids
increases solid solubility
Some have larger changes in solubility than others
Some solids even decrease in solubility

13. Heats of Solution
Dissolution of some ionic solids release energy and others absorb energy
How would each feel on the outside of beaker?
Which is endothermic/exothermic?
Solute particles must be separated
Solvent particles must be separated
Solvent and solute particles mix

14. Heats of Solution
Usually energy is required for the first two steps to occur
Amount depends on the strength of IMFs of each
Usually, energy is released in step three
Depends on how strong the IMFs are between the solute and solvent
Heat of solution
Net amount of heat absorbed or released when solute is dissolved in a solvent

15. Heats of Solution Heat of Solution = Step 1 + Step 2 + Step 3
Is negative (exothermic) when sum of the positive values from Steps 1&2 is less than the negative value in Step 3