﻿ Concept Summary - Solutions

# SolutionsConcept Summary

Nature of Solutions

· Solutions are homogeneous mixtures composed of two or more substances.

o They combine to form a single phase, generally the liquid phase.

o Solvent particles surround solute particles via electrostatic interactions in a process called solvation or dissolution.

o Aqueous solutions are most important for the MCAT; solvation in water can also be called hydration.

o Most dissolutions are endothermic, although the dissolution of gas into liquid is exothermic.

· Solubility is the maximum amount of a solute that can be dissolved in a given solvent at a given temperature; it is often expressed as molar solubility—the molarity of the solute at saturation.

· Complex ions or coordination compounds are composed of metallic ions bonded to various neutral compounds and anions, referred to as ligands.

o Formation of complex ions increases the solubility of otherwise insoluble ions (the opposite of the common ion effect).

o The process of forming a complex ion involves electron pair donors and electron pair acceptors such as those seen in coordinate covalent bonding.

Concentration

· There are many ways of expressing concentration.

o Percent composition by mass (mass of solute per mass of solution times 100%) is used for aqueous solutions and solid-in-solid solutions.

o The mole fraction (moles of solute per total moles) is used for calculating vapor pressure depression and partial pressures of gases in a system.

o Molarity (moles of solute per liters of solution) is the most common unit for concentration and is used for rate laws, the law of mass action, osmotic pressure, pH and pOH, and the Nernst equation.

o Molality (moles of solute per kilograms of solvent) is used for boiling point elevation and freezing point depression.

o Normality (number of equivalents per liters of solution) is the molarity of the species of interest and is used for acid—base and oxidation—reduction reactions.

Solution Equilibria

· Saturated solutions are in equilibrium at that particular temperature.

· The solubility product constant (Ksp) is simply the equilibrium constant for a dissociation reaction.

· Comparison of the ion product (IP) to Ksp determines the level of saturation and behavior of the solution:

o IP < Ksp: the solution is unsaturated, and if more solute is added, it will dissolve

o IP = Ksp: the solution is saturated (at equilibrium), and there will be no change in concentrations

o IP > Ksp: the solution is supersaturated, and a precipitate will form

· Formation of a complex ion in solution greatly increases solubility.

o The formation or stability constant (Kf) is the equilibrium constant for complex formation. Its value is usually much greater than Ksp.

o The formation of a complex increases the solubility of other salts containing the same ions because it uses up the products of those dissolution reactions, shifting the equilibrium to the right (the opposite of the common ion effect).

o The common ion effect decreases the solubility of a compound in a solution that already contains one of the ions in the compound. The presence of that ion in solution shifts the dissolution reaction to the left, decreasing its dissociation.

Colligative Properties

· Colligative properties are physical properties of solutions that depend on the concentration of dissolved particles but not on their chemical identity.

· Vapor pressure depression follows Raoult’s law.

o The presence of other solutes decreases the evaporation rate of a solvent without affecting its condensation rate, thus decreasing its vapor pressure.

o Vapor pressure depression also explains boiling point elevation—as the vapor pressure decreases, the temperature (energy) required to boil the liquid must be raised.

· Freezing point depression and boiling point elevation are shifts in the phase equilibria dependent on the molality of the solution.

· Osmotic pressure is primarily dependent on the molarity of the solution.

· For solutes that dissociate, the van't Hoff factor (i) is used in freezing point depression, boiling point elevation, and osmotic pressure calculations.

﻿