The Solution Process

• A solution is formed when one substance dispersed uniformly throughout another. The ability of substances to form solutions depends on two factors: (1) the natural tendency of substances to mix and spread into larger volumes when not restrained in some way and (2) the types of intermolecular interactions involved in the solution process.

The Natural Tendency toward Mixing

Entropy

• When the molecules mix and become more randomly distributed, there is an increase in a thermodynamic quantity.

Spontaneous process

• The mixing of gases.

• It occurs of its own accord without any input of energy from outside the system.

• Increases the entropy of the system.

Part of the enthalpy of the system

• All of the intermolecular forces (dispersion forces, dipole–dipole interactions, hydrogen bonding) and their corresponding energies.

• Decreasing the enthalpy of the system would correspond to increasing favorable intermolecular interactions between gas particles.

The Effect of Intermolecular Forces on Solution Formation

Three kinds of intermolecular interactions are involved in solution formation:

1. Solute–solute interactions between solute particles must be overcome to disperse the solute particles through the solvent.

2. Solvent–solvent interactions between solvent particles must be overcome to make room for the solute particles in the solvent.

3. Solvent–solute interactions between the solvent and solute particles occur as the particles mix.

• The extent to which one substance is able to dissolve in another depends on the relative magnitudes of these three types of interactions. Solutions form when the magnitudes of the solvent–solute interactions are either comparable to or greater than the solute–solute and solvent–solvent interactions.

Solvation

• Interactions such as this between solute and solvent molecules.

• When the solvent is water, the interactions are referred to as hydration.

Energetics of Solution Formation

• Solution processes are typically accompanied by changes in enthalpy.

Hydrates

• Compounds such as NiCl2 ‘ 6 H2O(s) that have a defined number of water molecules in the crystal lattice.

Saturated Solutions and Solubility

Crystallization

• This process, which is the opposite of the solution process.

• When the rates of these opposing processes become equal, a dynamic equilibrium is established, and there is no further increase in the amount of solute in solution.

Saturated

• A solution that is in equilibrium with undissolved solute.

Solubility

• The amount of solute needed to form a saturated solution in a given quantity of solvent.

• The solubility of a particular solute in a particular solvent is the maximum amount of the solute that can dissolve in a given amount of the solvent at a specified temperature, assuming that excess solute is present.

Unsaturated

• If we dissolve less solute than the amount needed to form a saturated solution.

Supersaturated

• Under suitable conditions, it is possible to form solutions that contain a greater amountof solute than needed to form a saturated solution.

Factors Affecting Solubility

Solute–Solvent Interactions

• When other factors are comparable, the stronger the attractions between solute and solvent molecules, the greater the solubility of the solute in that solvent.

• Because of favorable dipole–dipole attractions between polar solvent molecules and polar solute molecules, polar liquids tend to dissolve in polar solvents.

• Water is both polar and able to form hydrogen bonds.Thus, polar molecules, especially those that can form hydrogen bonds with water molecules, tend to be soluble in water.

Miscible

• Liquids that mix in all proportions, such as acetone and water.

Immiscible

• Do not dissolve in one another.

• Hexane, a hydrocarbon, is immiscible with water. Hexane is the top layer because it is less dense than water.

Alcohols

• Organic compounds with this molecular feature.

Over years of study, examination of different solvent–solute combinations has led to an important generalization:

Substances with similar intermolecular attractive forces tend to be soluble in one another.

Pressure Effects

• The solubilities of solids and liquids are not appreciably affected by pressure, whereas the solubility of a gas in any solvent is increased as the partial pressure of the gas above the solvent increases.

• The relationship between pressure and gas solubility is expressed by Henry’s law:

• Sg = kPg

Henry’s law constant

• is the ratio of a compound's partial pressure in air to the concentration of the compound in water at a given temperature.

Temperature Effects

• The solubility of most solid solutes in water increases as the solution temperature increases.

• In contrast to solid solutes, the solubility of gases in water decreases with increasing temperature.

Expressing Solution Concentration

• The concentration of a solution can be expressed either qualitatively or quantitatively. The terms dilute and concentrated are used to describe a solution qualitatively. A solution with a relatively small concentration of solute is said to be dilute; one with a large concentration is said to be concentrated.

Mass Percentage, ppm, and ppb

• One of the simplest quantitative expressions of concentration is the mass percentage of a component in a solution, given by

• Mass % of component = = mass of component in soln/ total mass of soln x 100

• Because percent means “per hundred,”

• We often express the concentration of very dilute solutions in parts per million (ppm) or parts per billion (ppb).

Mole Fraction, Molarity, and Molality

• Concentration expressions are often based on the number of moles of one or more components of the solution. The mole fraction of a component of a solution is given by

• Mole fraction of component = moles of component/total moles of all components

• The symbol X is commonly used for mole fraction, with a subscript to indicate the component of interest.

• Mole fractions are very useful when dealing with gases, but have limited use when dealing with liquid solutions. The molarity (M) of a solute in a solution is defined as

• Molarity = moles of solute/liters of soln

• The molality of a solution, denoted m, is a concentration unit that is also based on moles of solute. Molality equals the number of moles of solute per kilogram of solvent:

• Molality = moles of solute/kilograms of solvent

The definitions of molarity and molality are similar enough that they can be easily confused.

• Molarity depends on the volume of solution.

• Molality depends on the mass of solvent.

Colligative Properties

Colligative properties

• Lowering of the freezing point and raising of the boiling point are physical properties of solutions that depend on the quantity (concentration) but not on the kind or identity of the solute particles.

• (Colligative means “depending on the collection”; colligative properties depend on the collective effect of the number of solute particles.)

Vapor–Pressure Lowering

Vapor pressure

• Is the pressure exerted by the vapor when it is at equilibrium with the liquid (that is, when the rate of vaporization equals the rate of condensation).

Nonvolatile

• A substance that has no measurable vapor pressure.

Volatile

• Exhibits a vapor pressure.

Mixing a volatile liquid solvent and a nonvolatile solute increases entropy, forming a spontaneous solution. This stabilizes solvent molecules in liquid form, reducing their inclination to vaporize. Thus, a nonvolatile solute lowers the solvent's vapor pressure.

Raoult’s law

• States that a solvent's partial vapor pressure in a solution (or mixture) is equal or identical to the vapor pressure of the pure solvent multiplied by its mole fraction in the solution.

Ideal solution

• Defined as one that obeys Raoult’s law.

• Is a mixture in which the molecules of different species are distinguishable, however, unlike the ideal gas, the molecules in ideal solution exert forces on one another.

Boiling-Point Elevation

Van’t Hoff factor

• Defined as the ratio of the concentration of particles formed when a substance is dissolved to the concentration of the substance by mass.

• The molal boiling-point-elevation constant:

• ΔTb = i×Kb×m

Freezing-Point Depression

Molal freezing-point-depression constant

• The proportionality constant, Kf, is called the molal freezing-point depression constant.

Molal Boiling-Point-Elevation and Freezing-Point-Depression Constants

Osmosis

• The net movement of solvent is always toward the solution with the lower solvent (higher solute) concentration, as if the solutions were driven to attain equal concentrations.

• Osmosis plays an important role in living systems.

Osmotic pressure

• This pressure, which stops osmosis.

Crenation

• This causes the cell to shrivel.
  

Hemolysis

• Placing the cell in a solution that is hypotonic relative to the intracellular fluid causes water to move into the cell, which may cause the cell to rupture.

Colloids

Colloidal dispersions

• Is composed of solid, liquid, or gas particles dispersed in a continuous phase (solid, liquid, or gas).

• Colloids form the dividing line between solutions and heterogeneous mixtures. Like solutions, colloids can be gases, liquids, or solids.

Types of Colloids

Tyndall effect

• This scattering of light by colloidal particles.

Hydrophilic and Hydrophobic Colloids

Hydrophilic (“water loving”)

• In the human body, the extremely large protein molecules such as enzymes and antibodies

• Are kept in suspension by interaction with surrounding water molecules. A hydrophilic molecule folds in such a way that its hydrophobic groups are away from the water molecules, on the inside of the folded molecule, while its hydrophilic, polar groups are on the surface, interacting with the water molecules.

Hydrophobic (“water fearing”)

• Hydrophobic colloids can be dispersed in water only if they are stabilized in some way. Otherwise, their natural lack of affinity for water causes them to separate from the water.

Emulsify

• Means “to form an emulsion,” a suspension of one liquid in another.

Emulsifying agent

• A substance that aids in the formation of an emulsion.

Colloidal Motion in Liquids

Brownian motion

• The random zig-zag motion of a particle that is usually observed under high power ultramicroscope.

In 1905, Einstein developed an equation for the average square of the displacement of a colloidal particle, a historically very important development.