Solutions: Types and Characteristics

• Solutions are homogenous mixtures of two or more components, where the solvent is the component present in the largest quantity.
• Solvents determine the physical state of the solution, while solutes are the other components present.

Types of solutions based on solute and solvent nature:

• Gaseous solutions can be gas in gas (oxygen and nitrogen mixture), liquid in gas (chloroform in nitrogen), or solid in gas (camphor in nitrogen).
• Liquid solutions include gas in liquid (oxygen in water), liquid in liquid (ethanol in water), and solid in liquid (glucose in water).
• Solid solutions consist of gas in solid (hydrogen in palladium), liquid in solid (mercury in sodium amalgam), and solid in solid (copper in gold).

Types of solutions based on solute dissolution in solvent:

• Unsaturated solutions can dissolve more solute at a given temperature.
• Saturated solutions cannot dissolve more solute at a given temperature.
• Supersaturated solutions contain more solute than a saturated solution at a specific temperature; excess solute crystallizes quickly if the temperature is reduced.

Types of solutions based on the solvent:

• Aqueous solutions use water as the solvent, like salt, sugar, or copper sulfate in water.
• Non-aqueous solutions use a solvent other than water, for example, iodine in carbon tetrachloride.

Types of solutions based on solute and solvent amount:

• Dilute solutions contain a small amount of solute in a large amount of solvent.
• Concentrated solutions contain a large amount of solute in a small amount.

Types of solutions based on component uniformity:

• Homogenous solutions have uniform composition and properties throughout, like a cup of coffee.
• Heterogeneous solutions have non-uniform composition and properties, like oil and water.

Quantitative ways to expressing concentration of solutions include:

• Mass percentage (m/m) is the mass of a component in the solution divided by the total mass of the solution, multiplied by 100.
• Volume percentage (V/V) is the volume of a component in the solution divided by the total volume of the solution, multiplied by 100.
• Mass by volume percentage is the mass of solute dissolved in 100 mL of solution.
• Parts per million (ppm) indicates parts of one substance in one million parts of another, whether by mass or volume.
• ppm is calculated as (mass of solute / mass of solvent) x 1,000,000 or (volume of solute / volume of solvent) x 1,000,000, ensuring the same units for both solute and solvent.
• Mole fraction is the number of moles of a component divided by the total number of moles of all components, useful in relating vapor pressure to concentration.
• In a binary mixture, the mole fraction of A (XA) = nA / (nA + nB), where nA and nB are the moles of A and B.
• Molarity (M) is the moles of solute per liter of solution, measured in mol/L.
• Molality (m) is the moles of solute per kilogram of solvent, expressed in mol/Kg.

Solubility:

• Solubility is the maximum amount of a substance that can dissolve in a specified amount of solvent at a specified temperature, dependent on solute and solvent nature, temperature, and pressure.
• Solids dissolve as polar solutes in polar solvents and nonpolar solutes in nonpolar solvents.
• Whether a solute dissolves in a solvent depends on similar intermolecular interactions.
• Temperature significantly affects the solubility of a solid in a liquid.
• If the dissolution process is endothermic (Δsol H > 0) in a nearly saturated solution, solubility increases with temperature; if exothermic (Δsol H < 0), solubility decreases.
• Pressure has virtually no effect on solids in liquids due to their incompressibility.

Solubility of a Gas in Liquid:

• The extent to which many gasses dissolve in water varies; hydrogen chloride gas (HCl) is highly soluble, but oxygen dissolves only to a small extent.
• Gas solubility in liquids is affected by pressure: solubility increases with pressure, as described by Henry's law.
• Henry's law states that at constant temperature, the solubility of a gas is directly proportional to the partial pressure of the gas above the solution.
• The partial pressure of gas in the vapor phase (p) is proportional to the mole fraction of the gas (x) in the solution: p = KHx, where KH is Henry's law constant.
• Henry's law constant (KH) varies for different gases and is temperature-dependent.
• Henry's law applications include carbonating soft drinks under high CO2 pressure.
• Scuba divers face high dissolved gas concentrations due to increased pressure underwater.
• Rapid ascent causes nitrogen bubbles in the blood, leading to "bends" (painful/dangerous).
• At high altitudes, low oxygen partial pressure can cause anoxia (weakness, unclear thinking) due to low oxygen concentrations in blood/tissues.

Raoult's Law:

• Raoult's law states for any solution the partial vapor pressure of each volatile component in the solution is directly proportional to its mole fraction.
• For a binary solution with a non-volatile solute, only solvent molecules contribute to vapor pressure.
• If p1 is the solvent vapor pressure, x1 is the mole fraction, and p1º is the vapor pressure in the pure state, then according to Raoult's law p1 = x1p°1.
• The proportionality constant is equal to the vapor pressure of pure solvent or p°1.

Ideal and Non-Ideal Solutions:

• Ideal solutions obey Raoult's law over the entire concentration range.
• Enthalpy and volume of mixing are both zero (ΔHmix = 0, ΔVmix = 0); no heat is absorbed/evolved.
• Non-ideal solutions do not obey Raoult's law and exhibit vapor pressures higher (positive deviation) or lower (negative deviation) than predicted.
• Azeotropes are binary mixtures with the same composition in liquid and vapor phases, boiling at a constant temperature, making separation by fractional distillation impossible.
• Positive deviation from Raoult's law results in minimum boiling azeotropes.
• Azeotrope composition gives a solution containing approximately 95% by volume of ethanol.
• Negative deviation from Raoult's law results in maximum boiling azeotropes, such as nitric acid and water with ~68% nitric acid.

Colligative Properties:

• Colligative properties depend on the number of solute particles, regardless of their nature, relative to the total particles in the solution.
• Colligative properties include: relative lowering of vapor pressure, depression of freezing point, elevation of boiling point, and osmotic pressure.
• Vapor pressure lowering depends only on solute particle concentration, Raoult established.
• p1 = x1 pº1 establishes a relation between solution vapor pressure, mole fraction, and solvent vapor pressure.
• In solutions with non-volatile solutes, vapor pressure lowering depends on the sum of the mole fractions of different solutes.
• The expression on the left hand side of the equation is called relative lowering of vapor pressure and is equal to the mole fraction of the solute.
• Here n1 and n2 are the number of moles of solvent and solute respectively present in the solution.
• For dilute solutions n2 << n1, hence neglecting n2 in the denominator we have.

Elevation of boiling point:

• Elevation of boiling point depends on the number of solute molecules.
• The increase in the boiling point ∆Tb = ∆Tb -∆Tb (boiling point of solution - the boiling point of pure solvent) is known as the elevation of the boiling point.
• The molal elevation constant is also known as the Ebullioscopic Constant, with a unit of K kg mol-1.

Depression of freezing point:

• A solution will freeze when its vapor pressure equals the vapor pressure of the pure solid solvent
• Adding a non-volatile solid to a solvent decreases vapor pressure, lowering the freezing point of the solvent.
• The Freezing Point Depression Constant has a unit of K kg mol-1.

Osmotic pressure:

• Semipermeable membranes (SPM) allow small solvent molecules to pass, but hinder larger solute molecules.
• Osmotic pressure is the excess pressure required to prevent osmosis.
• Osmotic pressure (Π) is proportional to molarity (C) at temperature (T): P = CRT, where R is the gas constant.
• Measurement of osmotic pressure provides another method of determining molar masses of solutes.

Abnormal Molar Mass:

• Ionic compounds dissociate into ions when dissolved in water.
• The van't Hoff factor (i) accounts for dissociation or association extent.
• The van't Hoff factor (i) is defined as the ratio of normal molar mass to observed molar mass.
• The factor i is also defined as the actual number of particles in solution after dissociation to the number of formula units initially dissolved in solution.
• For association, i is less than 1; for dissociation, i is greater than 1.