# 13.S: Properties of Solutions (Summary)

These is a summary of key concepts of the chapter in the Textmap created for "Chemistry: The Central Science" by Brown et al.

## 13.1: The Solution Process

• interaction between solute and solvent molecules
• hydration – solvation when solvent is water

13.1.1 Energy Changes and Solution Formation

• overall enthalpy change in formation of a solution
• $$\Delta H_{\mathit{soln}}=\Delta H_1+\Delta H_2+\Delta H_3$$
• $$\Delta H_1=$$ separation of solute molecules
• $$\Delta H_1=$$ separation of solvent molecules
• $$\Delta H_3=$$ formation of solute-solvent interactions
• separation of solute particles is endothermic
• separation of solvent is endothermic
• third is exothermic
• formation of solution can be either exothermic or endothermic
• exothermic processes are spontaneous
• solution will not form if enthalpy is too endothermic
• H3 has to be comparable to H1 + H2
• Ionic substances cannot dissolve in nonpolar liquids
• Polar liquids do not form solutions with nonpolar liquids

13.1.2 Solution Formation, Spontaneity, and Disorder

• two nonpolar substances dissolve in one another
• attractive forces = London dispersion forces
• two factors in processes that are spontaneous: energy and disorder
• processes in which the energy content of the system decreases tend to occur spontaneously
• exothermic
• processes in which the disorder of the system increases tend to occur spontaneously
• solutions will form unless solute-solute or solvent-solvent interactions too strong relative to solute-solvent interactions

13.1.3 Solution Formation and Chemical Reactions

• distinguish between physical process of solution formation from chemical process that leads to a solution

## 13.2: Saturated Solutions and Solubility

• crystallization – reverse process of solution
• dynamic equilibrium – when equilibrium exists between process of solution and crystallization
• solute said to be saturated
• solubility – amount of solute needed to saturate a solution
• unsaturated – when there isn’t enough solute to saturate a solution
• supersaturated – when there is more solute than needed to saturate a solution
• for most salts crystallization of excess solute is exothermic

## 13.3: Factors Affecting Solubility

Solute-Solvent Interactions

• solubility increases with increasing molar mass
• London dispersion forces increase with increasing size and mass of gas molecules
• Miscible – pairs of liquids that mix in all proportions
• Immiscible – opposite of miscible
• Hydrogen-bonding interactions between solute and solvent leads to high solubility
• Substances with similar intermolecular attractive forces tend to be soluble in one another
• "like dissolves like"

Pressure Effects

• solubility of a gas in any solvent increases as pressure of gas over solvent increases
• relationship between pressure and solubility: Henry’s Law:
• $$C_g = kP_g$$
• Cg solubility of gas in solution phase (usually expressed as molarity), Pg partial pressure of gas over solution, k is proportionality constant (Henry’s Law constant)
• Henry’s law constant different fore each solute-solvent pair, and temperature

Temperature Effects

• solubility of most solid solutes in water increases as temperature of solution increase
• solubility of gases in water decreases with increasing temperature
• decreases solubility of O2 in water as temperature increases in one the effects of thermal pollution

## 13.4: Ways of Expressing Concentration

• dilute and concentrated used to describe solution qualitatively
• mass percentage of component in solution:

$\displaystyle\textit{mass% of component} = \frac{\textit{mass of component in soln}}{\textit{total mass of soln}}\times 100$

• very dilute solutions expressed in parts per million (ppm)

$\displaystyle\textit{ppm of component} = \frac{\textit{mass of component in soln}}{\textit{total mass of soln}}\times 10^6$

• 1ppm = 1g solute for each (106) grams of solution or 1mg solute per kg solution
• 1ppm = 1mg solute/L solution
• 1 ppb = 1g of solute/109 grams of solution or 1 m g solute/ L of solution

13.4.1 Mole Fraction, Molarity, and Molality

$\displaystyle\textit{mole fraction of component} = \frac{\textit{moles of component}}{\textit{total moles of all components}}$

• sum of mole fractions of all components of solution must equal one
• $$\displaystyle\textit{molarity} = \frac{\textit{moles solute}}{\textit{liters soln}}$$
• $$\displaystyle\textit{molality} = \frac{\textit{moles solute}}{\textit{kilograms of solvent}}$$
• molality goes not vary with temperature
• molarity changes with temperature because of expansion and contraction of solution

## 13.5: Colligative Properties

Colligative properties are physical properties that depend on quantity

Lowering the Vapor Pressure

• vapor pressure over pure solvent higher than that over solution
• vapor pressure needed to obtain equilibrium of pure solvent higher than that of solution

Raoult’s Law

• Raoult’s law: $$P_A = X_AP_{A}^°$$
• PA = vapor pressure of solution, XA = mole fraction of solvent, PA° = vapor pressure of the pure solvent
• Ideal solution – solution that obeys Raoult’s law
• Solute concentration is low, solute and solvent have similar molecular sizes and similar types of intermolecular attractions

Boiling-Point Elevation

• normal boiling point of pure liquid is the temperature at which pressure is 1 atm
• addition of a nonvolatile solute lowers vapor pressure of solution
• $$\Delta T_b=K_b m$$
• Kb = molal boiling-point-elevation constant
• Depends only on solvent
• boiling point elevation proportional to number of solute particles present in given quantity of solution

Freezing-Point Depression

• freezing point of solution is temperature at which the first crystals of pure solvent form in equilibrium
• freezing point of solution lower than pure liquid
• freezing point directly proportional to the molality of the solute:
• $$\Delta T_f=K_f m$$
• Kf = molal freezing-point-depression constant

Osmosis

• semipermeable – membranes that allow passage of some molecules and not others
• osmosis – the net movement of solvent molecules from the less concentrated solution to the more concentrated solution
• net movement of solvent always toward the solution with the higher solute concentration
• osmotic pressure – pressure needed to prevent osmosis, p
• $$\pi = \left(\frac{n}{V}\right)RT=MRT$$
• M = molarity of solution
• if solutions identical osmosis will not occur and said to be isotonic
• if one solution lower osmotic pressure = hypotonic, the solution that has higher osmotic pressure = hypertonic
• crenation = when cells shrivel up from the loss of water
• hemolysis = when cells rupture due to to much water

Determination of Molar Mass

• colligative properties can be used to find molar mass

## 13.6: Colloids

Colloidal dispersions (colloids) are intermediate types of dispersions or suspensions

• intermediate solutions between solutions and heterogeneous mixtures
• colloids can be gases, liquids, or solids
• colloid particles have size between 10 - 2000Å
• tyndall effect – scattering of light by colloids

Hydrophilic and Hydrophobic Colloids

• hydrophilic – colloids in which the dispersion medium is water
• hydrophobic – colloids not dispersed in water
• hydrophobic colloids have to be stabilized before being put in water
• natural lack of affinity for water causes separation
• can be stabilized by the adsorption of ions on the surface
• adsorped ions interact with water
• can also be stabilized by presence of other hydrophilic groups on surface

Removal of Colloidal Particles

• coagulation – enlarging colloidal particles
• heating or adding an electrolyte to mixture
• heating increases number of collisions and particles stick together increasing their size
• electrolytes causes neutralization of the surface charges of the particles which remove the electrostatic repulsion
• dialysis – use of semipermeable membranes to filter out colloidal particles