Chapter 9: Solutions – Study Notes for GOB Chemistry

Chapter 9: Solutions

9.1 Mixtures and Solutions

Mixtures are combinations of two or more substances that are not chemically bonded. They can be classified based on the uniformity of their composition.

• Heterogeneous mixtures: Non-uniform composition; different regions have different properties. Example: Rocky Road ice cream.

• Homogeneous mixtures: Uniform composition throughout. Example: Seawater.

Homogeneous mixtures are further classified as:

• Solutions: Contain particles the size of typical ions or small molecules (generally <2 nm).

• Colloids: Contain particles ranging from 2 to 500 nm in diameter.

Parts of a solution:

• Solute: The substance dissolved in the solvent (minor component).

• Solvent: The substance in which the solute is dissolved (major component).

Type	Particle Size	Appearance	Separation by Filtration
Solution	<2 nm	Clear	No
Colloid	2–500 nm	Cloudy	No
Heterogeneous Mixture	>500 nm	Cloudy/settles	Yes

Additional info: Table inferred from context as "Some Characteristics of Solutions, Colloids, and Heterogeneous Mixtures".

9.2 The Solution Process

The ability of a solute to dissolve in a solvent depends on the relative strength of intermolecular forces between solute and solvent compared to those within the pure substances.

• "Like dissolves like": Polar solvents dissolve polar and ionic solutes; nonpolar solvents dissolve nonpolar solutes.

• Example: Water (polar) dissolves salt (ionic), but not oil (nonpolar).

• Solvation (hydration): Water molecules surround and stabilize ions in solution.

• Dissolution is a physical change—the chemical identities of solute and solvent remain unchanged.

• Enthalpy change (ΔH): Dissolution can be exothermic (releases heat) or endothermic (absorbs heat).

Example: Hexane (nonpolar) and carbon tetrachloride (nonpolar) form a solution; octane (nonpolar) and methyl alcohol (polar) do not.

9.3 Solubility

Solubility describes how much solute can dissolve in a given amount of solvent at a specific temperature.

• Miscible: Two liquids that are soluble in all proportions (e.g., ethanol and water).

• Saturated solution: Contains the maximum amount of dissolved solute at equilibrium; excess solute remains undissolved.

• Solubility: Maximum amount of solute that dissolves in a given solvent at a specified temperature (e.g., 35.8 g NaCl/100 mL water at 20°C).

9.4 The Effect of Temperature on Solubility

Temperature affects the solubility of solids and gases differently.

• Solids: Most become more soluble as temperature increases.

• Gases: Become less soluble as temperature increases.

• Supersaturated solution: Contains more solute than a saturated solution at a given temperature; unstable and precipitates when disturbed.

Example: The solubility of oxygen in water decreases from 8.3 mg/L at 25°C to 7.0 mg/L at 35°C.

9.5 The Effect of Pressure on Solubility: Henry’s Law

Henry’s Law describes the relationship between the solubility of a gas and its partial pressure above the solution.

• Statement: The solubility of a gas is directly proportional to its partial pressure at constant temperature.

• Equation:

$C = kP$

• C: Concentration (solubility) of the gas

• k: Henry’s law constant (depends on gas and solvent)

• P: Partial pressure of the gas

For two conditions at constant temperature:

$\frac{C_1}{P_1} = \frac{C_2}{P_2}$

• Le Châtelier’s Principle: Increasing pressure increases gas solubility; decreasing pressure decreases it.

Example: The solubility of oxygen in blood decreases at high altitudes due to lower atmospheric pressure.

9.6 Units of Concentration

Concentration expresses the amount of solute in a given amount of solution. Several units are commonly used:

Unit	Definition	Formula
Mass/mass % (m/m)%	g solute/100 g solution	$\text{(m/m)%} = \frac{\text{g solute}}{\text{g solution}} \times 100$
Volume/volume % (v/v)%	mL solute/100 mL solution	$\text{(v/v)%} = \frac{\text{mL solute}}{\text{mL solution}} \times 100$
Mass/volume % (m/v)%	g solute/100 mL solution	$\text{(m/v)%} = \frac{\text{g solute}}{\text{mL solution}} \times 100$
ppm	parts per million	$\text{ppm} = \frac{\text{g solute}}{\text{g solution}} \times 10^6$
ppb	parts per billion	$\text{ppb} = \frac{\text{g solute}}{\text{g solution}} \times 10^9$
Molarity (M)	mol solute/L solution	$M = \frac{\text{mol solute}}{\text{L solution}}$

• Example (m/m)%: 18-carat gold contains 75% gold by mass.

• Example (v/v)%: To make 75 mL of a 5.0% (v/v) methyl alcohol solution, use 3.75 mL methyl alcohol.

• Example (m/v)%: 1.8 g heparin sodium in 15 mL solution is 12% (m/v).

• Example (ppm): 100 ppb chloroform in 400 g water allows 0.00004 g chloroform.

• Example (Molarity): Dissolving 2.355 g H2SO4 in 50.0 mL yields 0.480 M solution.

9.7 Dilution

Solutions are often prepared by diluting a more concentrated stock solution. The amount of solute remains constant during dilution; only the volume changes.

The dilution equation:

$M_1 V_1 = M_2 V_2$

• M1: Initial molarity

• V1: Initial volume

• M2: Final molarity

• V2: Final volume

Example: Diluting 75 mL of 3.5 M glucose to 450 mL yields a final concentration of 0.58 M.

9.8 Ions in Solution: Electrolytes

Ionic compounds in aqueous solution can conduct electricity due to the movement of ions.

• Strong electrolytes: Ionize completely in water (e.g., NaCl).

• Weak electrolytes: Partially ionize in water (e.g., acetic acid).

• Nonelectrolytes: Do not produce ions in water (e.g., sugar).

Electrolyte concentration is often expressed in equivalents (Eq):

• 1 Eq = 1 mol of charge

• 1 mEq = 0.001 Eq

Example: 5.0 mEq/L K+ in blood corresponds to a specific mass of potassium per liter.

9.9 Properties of Solutions (Colligative Properties)

Colligative properties depend on the number of dissolved particles, not their identity.

• Vapor pressure lowering: Solutions have lower vapor pressure than pure solvents.

• Boiling point elevation: Solutions boil at higher temperatures than pure solvents.

• Freezing point depression: Solutions freeze at lower temperatures than pure solvents.

• Osmosis: Solvent passes through a semipermeable membrane toward higher solute concentration.

Boiling point elevation equation:

$\Delta T_b = K_b m$

• Kb: Boiling point elevation constant (for water, 0.51°C·kg/mol)

• m: Molality (mol solute/kg solvent)

Freezing point depression equation:

$\Delta T_f = K_f m$

• Kf: Freezing point depression constant (for water, 1.86°C·kg/mol)

Example: 0.75 mol KBr in 1.0 kg water raises boiling point by 0.77°C (to 100.77°C).

9.10 Osmosis and Osmotic Pressure

Osmosis is the movement of solvent through a semipermeable membrane from a region of lower solute concentration to higher solute concentration.

• Osmotic pressure (π): The pressure required to prevent osmosis.

• Osmolarity (osmol/L): The total molarity of all solute particles in solution.

Osmotic pressure equation:

$\pi = MRT$

• M: Molarity of solute particles

• R: Gas constant

• T: Temperature in Kelvin

Isotonic: Same osmolarity as blood plasma (0.30 osmol/L). Hypotonic: Lower osmolarity than plasma. Hypertonic: Higher osmolarity than plasma.

Example: 5.0% (m/v) glucose solution has an osmolarity equal to its molarity because glucose does not dissociate.

9.11 Dialysis

Dialysis is the process by which small solute particles and solvent molecules pass through a semipermeable membrane, but larger particles (like proteins) do not.

• Hemodialysis: Used to remove waste from the blood of patients with kidney failure.

• Proteins: Do not cross the membrane; help maintain osmotic balance in blood.

• Filtration vs. Reabsorption: Blood pressure pushes water out of plasma (filtration); osmotic pressure of proteins draws water back in (reabsorption).

Application: Delivery of nutrients and removal of waste in the body is regulated by osmosis and dialysis.