Chapter 13 review - Lecture notes 13 PDF

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Chapter 13. Solutions 13.1 Types of Solutions: Intermolecular Forces and Solubility—a solute is that which dissolves in a solvent. A solvent is typically the more abundant component of the solution. The solubility is the maximum amount of solute that will dissolve in a fixed quantity of solvent (or solution) at a given temperature. Solubilities may be understood in terms of the relative strengths of the intermolecular forces within both the solute and the solvent and between them. A rule of thumb for solubility is “like dissolves like” which implies that substances with similar types of intermolecular forces dissolve in one another. Intermolecular forces may be identified based on chemical formulas, structures, sizes, shapes, polarities, and polarizabilities. 13.3 Why Substances Dissolve: Breaking Down the Dissolution Process The enthalpy of dissolution may be broken down into three key contributions: (1) the enthalpy of separating solute particles from each other (endothermic); (2) the enthalpy of separating solvent particles from each other (endothermic); and (3) the enthalpy of mixing the separated particles (exothermic). ° ° ° ° ∆ H soln=∆ H 2 + ∆ H 1 + ∆ H mix ° ° ∆ H °aq=∆ H lattice + ∆ H hydration

The charge densities of ions are important in determining both the lattice enthalpies and the ° ° hydration enthalpies of salts. Enthalpies of solution ( ∆ H soln ∧∆ H aq ) can be positive, negative, or near zero depending on the relative values of the endothermic ( ° ° ° ° ° ∆ H 2 , ∆ H 1 ,∧∆ H lattice ) and exothermic ( ∆ H mix ∧∆ H hydration ) contributions in the above equations. A positive entropy of dissolution can promote the solubility of a solute whose enthalpy of solution is thermodynamically unfavorable (endothermic). 13.4 Solubility as an Equilibrium Process A saturated solution is one that contains the maximum amount of dissolved solute at the temperature and pressure of the mixture. An unsaturated solution contains less than this amount and a supersaturated solution contains more than this amount, but is thermodynamically unstable. Solids are typically more soluble at higher temperatures, gases are typically less soluble at higher temperatures. Henry’s law is an equation that shows the dependence of gas solubility on the partial pressure of the gas over the solution. S gas=k H , gas P gas

Henry’s law also gives the partial pressure of a gas over a solution of a given concentration.

13.5 Concentration Calculations—given the masses of solute and solvent and the density of the solution, be able to: 1. Find the mole fractions of solute and solvent in the solution: n x 2= 2 ; a. n1 + n2 n1 b. x 1= . n1 + n 2 2. Find the molality of the solution (moles of solute per kilogram of solvent): n 1000. g molality= 2 × a. 1kg m1 3. Find the volume of the solution sample in liters: 1L 100. g × V total (L)= g a. 1000 mL ) ρ( mL 4. Find the Molarity of the solution (moles of solute per liter of solution): n2 Molarity= a. V total ( L ) 5. Calculate mass fraction: m2 m2 mass fraction= = . m 1 + m 2 1000. g+ m 2 Dilution calculations:

M 1 V 1= M 2 V 2

13.6 Colligative Properties of Solutions There are 4 colligative properties of solutions that depend on the number of dissolved solute particles, but not on their nature: Vapor pressure lowering:

∆ P1=x 2 ∆ P°1

∆ P1=i x2 ∆ P°1

Boiling point elevation:

∆ T b=K b m

∆ T b=i K b m

Freezing point depression:

∆ T f =K f m

∆ T f =i K f m

Osmotic pressure:

Π =CRT

Π =i CRT

Because strong electrolytes dissociate into component ions, they produce a stronger colligative effect due to the greater number of dissolved particles. This effect is taken into account in the second equation for each property in which the van’t Hoff i-factor is approximately equal to the number of ions produced per formula unit....