11.1 - Solution Composition

• A solute is a substance being dissolved. The solvent is the dissolving medium.

• Molarity= Moles of solute/ liters of solution

• Mole fraction: the ratio of the number of moles of a given component to the total number of moles of solution

• When liquids are mixed, the liquid present in the largest amount is called the solvent

• In very dilute aqueous solutions, the magnitude of the molality and the molarity are almost the same

• Since molarity depends on the volume of the solution, it changes slightly with temperature

• Molarity is independent of temperature because it depends only on mass

• The definition of an equivalent depends on the reaction taking place in the solution

• For oxidation-reduction reactions, the equivalent is defined as the quantity of oxidizing or reducing agents that can accept or furnish 1 mole of electrons

11.2 - The Energies of Solution Formation

• The solubility of various vitamins is important in determining correct dosages

• The insolubility of barium sulfate means it can be used safely to improve X rays of the gastrointestinal tract

• Polar solvents dissolve polar solutes; non-polar solvents dissolve non-polar solutes

• Steps to the formation of a liquid solution: Separating the solute into its individual components (expanding the solute).

  • Overcoming intermolecular forces in the solvent to make room for the solute

  • Allowing the solute and solvent to interact to form the solution.

• The enthalpy of solution is the sum of the energies used in expanding both solvent and solute and the energy of solvent-solute interaction

• The heat of hydration represents the enthalpy change associated with the dispersal of a gaseous solute in water.

• Processes naturally run in the direction that leads to the most probable state

• One factor that favors a process is an increase in the probability

• The heat of the solution is expected to be large and positive, and the large quantity of energy required acts to prevent the solution from forming

• Solution forms because of the increase in the probability of the mixed state

11.3 - Factors Affecting Solubility

• Vitamins can be divided into two classes: fat-soluble and water-soluble

  • The fat-soluble vitamins can build up in the fatty tissues of the body

  • If excessive amounts of these vitamins are consumed, their buildup can lead to the illness hypervitaminosis

  • In contrast, the water-soluble vitamins are excreted by the body and must be consumed regularly.

• While pressure has little effect on the solubilities of solids or liquids, it does significantly increase the solubility of a gas

  • Carbonated beverages, for example, are always bottled at high pressures of carbon dioxide to ensure a high concentration of carbon dioxide in the liquid.

• The relationship between gas pressure and the concentration of dissolved gas is given by Henry’s law: C=kP

  • Henry’s law states that the amount of a gas dissolved in a solution is directly proportional to the pressure of the gas above the solution

  • Henry’s law holds only when there is no chemical reaction between the solute and solvent.

• The dissolving of a solid occurs more rapidly at higher temperatures, but the amount of solid that can be dissolved may increase or decrease with increasing temperature

• The only sure way to determine the temperature dependence of a solid’s solubility is by experiment

• The boiler scale reduces the efficiency of heat transfer and can lead to blockage of pipes

11.4 - The Vapor Pressures of Solutions

• The pressure of vapor necessary to achieve equilibrium with the pure solvent is greater than that required to reach equilibrium with the aqueous acid solution

  • As the pure solvent emits vapor to attempt to reach equilibrium, the aqueous sulfuric acid solution absorbs vapor to try to lower the vapor pressure toward its equilibrium value

• A nonvolatile solute has no tendency to escape from solution into the vapor phase

  • The dissolved nonvolatile solute decreases the number of solvent molecules per unit volume and it should proportionately lower the escaping tendency of the solvent molecules

• Raoult’s law states that the vapor pressure of a solution is directly proportional to the mole fraction of solvent present

  • The phenomenon of the lowering of the vapor pressure gives us a convenient way to “count” molecules and thus provides a means for experimentally determining molar masses

  • Using Raoult’s law, we can determine the number of moles of solute present

  • Strong solute-solvent interaction gives a vapor pressure lower than that predicted by Raoult’s law

• The lowering of vapor pressure depends on the number of solute particles present in the solution

11.5 - Boiling- Point Elevation and Freezing- Point Depression

• Freezing-point depression, boiling-point elevation, and osmotic pressure are called colligative properties

  • They are grouped together because they depend only on the number, and not on the identity, of the solute particles in an ideal solution.

• Melting point and freezing point both refer to the temperature where the solid and liquid coexist

• The observed freezing-point depression can be used to determine molar masses and to characterize solutions.

• The normal boiling point of a liquid occurs at the temperature where the vapor pressure is equal to 1 atmosphere

• The liquid/vapor line is shifted to higher temperatures than those for pure water

• The presence of the solute lowers the rate at which molecules in the liquid return to the solid-state

• Because a solute lowers the freezing point of water, compounds such as sodium chloride and calcium chloride are often spread on streets and sidewalks to prevent ice from forming in freezing weather

11.6 - Osmotic Pressure

• A solution and pure solvent are separated by a semipermeable membrane, which allows solvent but not solute molecules to pass through

  • This flow of solvent into the solution through the semipermeable membrane is called osmosis

  • Osmosis can be prevented by applying pressure to the solution

  • The minimum pressure that stops the osmosis is equal to the osmotic pressure of the solution.

  • Osmotic pressure can be used to characterize solutions and determine molar masses, but osmotic pressure is particularly useful because a small concentration of solute produces a relatively large osmotic pressure

  • Measurements of osmotic pressure generally give much more accurate molar mass values than those from the freezing- point or boiling-point changes

  • In osmosis, a semipermeable membrane prevents the transfer of all solute particles

• If a solution in contact with pure solvent across a semipermeable membrane is subjected to an external pressure larger than its osmotic pressure, reverse osmosis occurs

• Various schemes have been suggested, including solar evaporation, reverse osmosis, and even a plan for towing icebergs from Antarctica. The problem, of course, is that all the available processes are expensive

11.7 - Colligative Properties of Electrolyte Solutions

• The colligative properties of solutions depend on the total concentration of solute particle

• At a given instant a small percentage of the sodium and chloride ions are paired and thus count as a single particle

• As the solution becomes more dilute, the ions are farther apart and less ion-pairing occurs.

• Ion pairing occurs to some extent in all electrolyte solutions

• The colligative properties of electrolyte solutions are described by including the van’t Hoff factor in the appropriate equation

• The Van’t Hoff factor I represents the number of ions produced by each formula unit of solute

11.8 - Colloids

• A suspension of tiny particles stabilized by electrostatic repulsion among the ion layers surrounding the individual particles

• The suspended particles are single large molecules or aggregates of molecules or ions ranging in size from 1 to 1000 nm

• It can be coagulated (destroyed) by heating or adding an electrolyte

• When a colloid is placed in an electric field, the dispersed particles all migrate to the same electrode and thus must all have the same charge.

• Since the colloidal particles all have an outer layer of ions with the same charge, they repel each other and do not easily aggregate to form particles that are large enough to precipitate