Understanding Solutions: Solutes, Solvents & Concentration

Solutions: Introduction

• Solutions consist of solute and solvent.
• Solute is the substance in smaller quantity.
• Solvent is the substance in larger quantity.
• In sugar water, sugar is the solute, and water is the solvent.
• Concentrated solutions have higher amounts of solute relative to solvent
• Diluted Solutions have lower amounts of solute relative to solvent.

Types of Solutions

• Solutions are categorized by their solvent.
• Aqueous solutions use water as the solvent.
• Non-aqueous solutions use a different solvent.
• Concentration is the amount of solute in a known amount of solvent.
• Concentration can be dilute or concentrated.

Expressing Concentration

• Gram per liter concentration is the amount of solute in grams per liter of solution.
  • 20 grams per liter of aqueous urea solution contains 20 grams of urea in 1 liter of water solution.
• Molarity (M) is moles of solute per liter of solution.
  • Molarity = (Given weight of solute / Molecular weight of solute) / Volume of solution in liters
  • For volume in milliliters: Molarity = (Given weight of solute / Molecular weight of solute) / (Volume of solution in mL) * 1000
  • Units for molarity: M, Mils Per KG, or mol/L.
• Molality (m) is moles of solute per kilogram of solvent.
  • molality = (Given weight of solute / Molecular weight of solute) / (weight of solvent in grams) * 1000

Mole Fraction

• Mole fraction (x) represents the ratio of moles of a component to the total moles in a solution.
• For components A and B: xA = moles of A / (moles of A + moles of B)
• The sum of all mole fractions in a solution equals 1.

Mass Percent

• Mass percent = (Mass of solute / Total mass of solution) * 100.

Solubility

• Solubility is the maximum amount of solute that dissolves in a solvent at a specific temperature.
• Factors affecting solubility: nature of solute and solvent, temperature, and pressure.

Solubility: Solid in Liquid

• "Like dissolves like" governs the nature of solute and solvent.
  • Polar solutes dissolve in polar solvents; non-polar solutes dissolve in non-polar solvents.
• Temperature affects solubility.
  • Exothermic reactions: solubility decreases with increasing temperature.
  • Endothermic reactions: solubility increases with increasing temperature.
• Pressure has an insignificant impact on solubility.

Solubility: Gas in Liquid

• Easily liquefied gases are more soluble.
• Gases that chemically combine with the solvent are more soluble.
• Temperature: Solubility of gases decreases with increasing temperature.
  • Gas molecules have greater kinetic energy and escape the solution.
• Pressure: Solubility of gases increases with increasing pressure; this is known as Henry's Law.

Henry’s Law

• The amount of gas that dissolves in a liquid is directly proportional to the gas's partial pressure above the liquid.
• Formula: p = KH * x, where p is the partial pressure, x is the mole fraction, and KH is Henry's constant.
• Henry's constant (KH) is inversely proportional to the gas's solubility.
• Applications:
  • Carbonated beverages are sealed under high pressure to increase carbon dioxide solubility.
  • At high altitudes, lower partial pressure of oxygen results in less oxygen dissolving in the blood, leading to hypoxia.

Vapor Pressure

• Vapor Pressure is the pressure exerted by vapors on the surface of a liquid at equilibrium.
• Vapors exert pressure on the liquid's surface when water boils.

Raoult's Law

• The partial vapor pressure of each component of an ideal mixture is equal to the vapor pressure of the pure component multiplied by its mole fraction in the mixture.
• Applicable to solutions with two volatile liquids (A and B)
• Formula: pA = p0A * xA
  • pA = partial vapor pressure of component A in the solution
  • p0A = vapor pressure of pure component A
  • xA = mole fraction of component A in the solution
• Dalton's Law: Total vapor pressure of a solution equals the sum of partial vapor pressures of the liquid.

Solutions: Non-Volatile Solute

• Non-volatile solutes do not contribute to vapor pressure.
  • pTotal = p0A * xA
• Non-volatile substances do not produce vapor.

Types of Solutions

• Ideal Solutions obey Raoult's Law.
• Non-ideal solutions do not obey Raoult's Law.

Characteristic of an Ideal Solution

• The magnitude difference between intermolecular forces separates ideal solutions from non-ideal ones.
  • Intermolecular forces hold the mixtures molecules in solution.

Non-Ideal Solutions

• Exhibit different interactions between molecules that differ in pure liquids.
• Positive Deviation: Intermolecular forces are weaker than in pure liquid.
• Negative Deviation: Intermolecular forces are stronger than in pure liquids.

Positive Deviation

• The vapor pressure of the solution is higher than predicted by Raoult’s Law.

Negative Deviation

• Intermolecular forces are stronger than in pure liquids, resulting in a total vapor pressure less than predicted.

Colligative Properties

• Properties depend on the number of solute particles, not their identity or qualities.
• Four main colligative properties:
  • Relative Lowering of Vapor Pressure.
  • Elevation in Boiling Point.
  • Depression in Freezing Point.
  • Osmotic Pressure.
• Colligative effects rise in proportion to the number of solute particles.

Relative Lowering of Vapor Pressure

• Adding a non-volatile solute lowers the vapor pressure of a solvent.
• Formula: (p0 - p) / p0 = xB
  • p0 = vapor pressure of the pure solvent
  • p = vapor pressure of the solution
  • xB = mole fraction of the solute

Elevation in Boiling Point

• A solution with a non-volatile solute has a higher boiling point than the pure solvent.
• Formula: ΔTb = Kb * m
  • ΔTb = elevation in boiling point
  • Kb = ebullioscopic constant (specific to the solvent)
  • m = molality of the solution

Depression in Freezing Point

• A solution with a non-volatile solute has a lower freezing point than the pure solvent.
• Formula: ΔTf = Kf * m
  • ΔTf = depression in freezing point
  • Kf = cryoscopic constant (specific to the solvent)
  • m = molality of the solution

Osmotic Pressure

• Pressure prevents solvent molecules from flowing across a semipermeable membrane from low to high solute concentration.
• Formula: π = cRT
  • π = osmotic pressure
  • c = molar concentration of the solution
  • R = ideal gas constant
  • T = absolute temperature

Azeotropes

• Azeotropes are mixtures of liquids that have constant boiling points and are non-separable via distillation.
• Form when the vapor and liquid compositions are equal.
• Two types of azeotropes:
  • Minimum boiling azeotrope: boiling point is less than that of individual components, exhibiting positive deviation.
  • Maximum boiling azeotrope: boiling point is higher than that of individual compounds and shows negative deviation from Raoult's Law.

Van't Hoff Factor

• The Van't Hoff Factor is the ratio of moles of particles in solution to moles of formula units dissolved.
  • It is a correction factor implemented to solutions that do not exhibit ideal behavior.
  • Formula: i = number of independent particles after solutions / number of moles before solutions

Non-Ideal Solution Types

• Isotonic Solutions: equal osmotic pressure with no movement through the membrane.
• Hypertonic Solutions: high concentration.
• Hypotonic Solutions: low concentration.
• Relationship between i and α is used to find association or dissociation, where α = Degree of dissociation/association.

Explore the basics of solutions, including the roles of solutes and solvents. Learn how to differentiate between aqueous and non-aqueous solutions. Understand concentration in solutions, including gram per liter concentration.