# Solubility, cont.

### Liquid solutions

#### "Like dissolves like."

• Solvation - the process of surrounding a solute with solvent molecules).
• Hydration (solvation with water).
• Saturation - the maximum amount of solute which can be in solution in equilibria with its pure state.
• Supersaturation - dissolving solute in excess of saturation - unstable to the addition of solute (carbonated beverages, honey, etc.) Video
• Temperature effects:
• Gases decrease in solubility with increasing temperature. [example of oxygen solubility]
• Many solids increase in solubility with increasing temperature, but not always true.
• Pressure effects:
• Gases increase in solubility with increasing pressure (Henry's Law: P = kC, where k is a constant specific to the solution). [example of carbonated beverages - p 519 in Zumdahl 9th).
• Note that Henry's Law only holds for gases that are non-reactive with the solvent - e.g. works for oxygen and nitrogen in water, does not work for HCl and water (goes to H+ + Cl-)
• Solubility of solids generally uneffected by pressure, since both liquids and solids essentially incompressible.

# Concentration Measures

## Concentration Terms

### Percent Concentration

• Mass percent
• ppt = parts/thousand (1g/L of water); ppm = parts/million (1mg/L of water); ppb = parts/billion (1 microgram/L of water)
• Volume percent

### Molarity = M = moles of solute dissolved in 1 L of solution.

The most commonly used concentration term in chemistry = moles of solute dissolved in 1 L of solution.

### FYI

Example: Make up a 1.00000 L solution of 0.25 M NaCl (note that water is the "default" solvent).

First weigh out o.25 moles of NaCl

= (0.25 mole)(22.99 g + 35.45 g)/mole = 14.61 g

Example: What is the concentration of a solution made by dissolving 10.00 g of KI in enough water to make
1.00000 L?

First need to find the number of moles of KI:

(10.00 g) / ({39.10 g+ 126.9 g}/mole) = 6.135 x 10-2 mole

Thus the concentration will be 6.135 x 10-2 M

### Mole fraction = X = moles of solute dissolved in total moles of solution, X = na /n

Example: What is the mole fraction of a solution of 10.0 moles of glycerol dissolved in 15.0 moles of water?

(10 mol) / (10 mol + 15 mol) = 10/25 = 0.400

# Colligative properties

Colligative properties (properties which depend only on the number or concentration, not on the type, of particles). [Exchange across surfaces model]

Colligative properties are only strictly followed for ideal solutions. That is, other forces are at work in real solutions, so will get deviations. As a result colligative properties are followed most closely for dilute solutions (e.g. <0.1 M) where solute-solute interactions are minimal.

• ### Vapor pressure lowering

Raoult's Law: P = XP°, where P = vapor pressure of substance in solution, P° = the vapor pressure of the pure substance and X = its mole fraction. Recall that mole fraction is the number of moles of substance divided by the total number of moles of all substances in the solution (moles solute/(moles solute + moles solvent)) In other words the vapor pressure of a substance in solution is proportional to the molecular fraction or molecular percentage of that substance in the solution.

Example: What is the vapor pressure of water in 80 proof alcohol at 25° C (vapor pressure = 23.76 mmHg).

P = XP°

XH2O = [60g/18.01 g/mol] / [60g/18.01 g/mol + 40g/(2x12.01 + 6x1.008 + 16.00)g/mol = 0.79

P = 0.79 (23.76 mmHg) = 18.77 mmHg = 19 mmHg

• ### Boiling point elevation

Tb = kbm, where m = molality = moles solute/kg solvent, and kb is a constant specific to the solvent.

Which of the following solutions will have the highest boiling point: 3 m glucose or 1 m aluminum chloride?

First need to look at concentration of particles.

glucose is covalent, so 1 m particles,

aluminum chloride is ionic with 1 mole aluminum ions and 3 moles of chloride ions for each mole of AlCl3 = 4 m particles,

Therefore the 1 m aluminum choride solution will have the higher bp.

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