Chapter 12: Solutions

Chemistry 1062: Principles of Chemistry II

Andy Aspaas, Instructor

Solutions

• Solution: homogeneous mixture of two or more substances (atoms, molecules, or ions)

• Can exist in any state of matter

Solute Solvent

Gaseous solution

Component in smaller amount

Component in larger amount

Liquid solutionLiquid component in smaller amount, or solid or gas

Liquid component in larger amount

Solid solution Component in smaller amount

Component in larger amount

States of solutions

• Miscible fluids: fluids that dissolve with each other in all proportions

• All nonreactive gases are generally miscible– Air is a gaseous solution (N2, O2, CO2, etc.)

• Liquid solutions: dissolving a solid, liquid, or gas into a liquid– Ethanol and water are miscible, and when mixed make a

liquid solution– Brine is solid sodium chloride dissolved in water

• Solid solutions are called alloys when metals are mixed– Dental fillings, brass, steel, etc.

Solubility and saturation

• If 40.0 g of NaCl were stirred in 100 mL of 20 °C water, most of the salt would dissolve but some would remain on the bottom

– Ions dissolve by leaving the surface of the crystal and entering the liquid solution

– Some crystals may re-deposit on the crystal

• Equilibrium is reached at point where particles dissolve at same rate as they return to the crystal

– The solution has become saturated

Solubility and saturation

• The point at which the solution becomes saturated can be expressed as solubility (at a given temperature)

– The solubility of NaCl in 20 °C water is 36.0 g NaCl / 100 mL H2O

• The solution is unsaturated when not enough solid has been added for the equilibrium to be reached

– Unsaturated solutions can support addition of more solid to be dissolved

Supersaturation

• Supersaturated solution: solution which contains more dissolved substance than a saturated solution does

• Sodium acetate and many other ionic compounds more soluble in hot water than in cold water

– If a saturated solution is prepared at high temperature, and then the temperature is slowly lowered, the solution may become supersaturated

– Introduction of any solid to a supersaturated solution will cause the whole solution to quickly crystallize

Factors behind solubility

• “Like dissolves like” - similar substances tend to dissolve in each other

• The natural tendency of substances to mix through the random motion of their particles can be overridden if one component has strong intermolecular forces, and the other does not

– Oil and water: water’s intermolecular forces are maintained if nonpolar oil molecules do not interrupt the water molecules

– Alcohol and water: similar hydrogen bonds can be formed, so 3-carbon alcohols and smaller are miscible in water (larger alcohols are too dissimilar to water)

Solubility of ionic compounds

• Partial charges of water molecules orient themselves towards oppositely charged ions in solutions (ion-dipole force)

• Hydration: water molecules surrounding ions - this favors dissolving of ionic solids in water

• Lattice energy: energy holding together ions in a crystal lattice - this works against dissolving– Lattice energy increases with ion charge– Lattice energy decreases with ionic radius– Hydration energy increases with ionic radius

Temperature effects on solubility

• Most gases are less soluble in water at higher temperatures (bubbles that appear when heating water)

• Most ionic solids are more soluble in water at higher temperatures

– Some have very little change, like NaCl– Some are less soluble in higher temperatures

• Heat of solution: heat absorbed or released when a solid is dissolved

– Depends on combination of lattice energy and hydration energy

– Chemical hot packs and cold packs take advantage of this

Pressure effects on solubility

• All gases become more soluble in a liquid at a given temperature when the partial pressure fo the gas over the solution is increased– Le Chatelier’s principle: if an equilibrium is

disturbed by a temperature, pressure or concentration change, the equilibrium will shift to compensate for the change

– Increasing the partial pressure of CO2(g) over water will cause more CO2 to dissolve

CO2(g) CO2(aq)• This equilibrium will shift to the right (more CO2 will

dissolve) to compensate for the pressure increase

Henry’s law

• Henry’s law: solubility of a gas is directly proportional to the partial pressure of the gas above the solution

S = kHPwhere S is solubility (mass of solute per unit

volume of solvent),

kH is Henry’s law constant (for gas and liquid at a given temperature),

P is partial pressure of the gas

Colligative properties and concentration

• Colligitave properties depend on concentration of solute molecules or ions in solution, but not the chemical identity of the solute

• Molarity, M = (moles solute)/(liters solution)

• Mass percentage: [(mass solute)/(mass solution)]*100%

• Molality, m = (moles solute)/(kilograms solvent)

• Mole fraction, XA = (moles A)/(total moles solution)

Vapor pressure of a solution

• Vapor pressure lowering: colligative property– Vapor pressure of pure solvent minus vapor

pressure of solution• Raoult’s law: PA = P°AXA

– if solute must be nonvolatile nonelectrolyte– PA = partial pressure of solvent– P°A = vapor pressure of pure solvent– XA = mole fraction of solvent in solution

• Or, ∆P = P°AXB where XB = mole fraction of solute

Distillation

• An ideal solution follows Raoult’s law for all mole fractions 0–1

– When two components are chemically similar, their intermolecular forces are similar

• Total vapor pressure, P = P°AXA + P°BXB

• Vapor over an ideal solution is richer in the more volatile component

• Fractional distillation condenses and re-vaporizes the solution many times to exploit this

Boiling point elevation

• Addition of a nonvolatile solute reduces the solvent’s vapor pressure

• Normal boiling point: temperature at which vapor pressure = 1 atm

• So, addition of solute requires a higher temperature in order for vapor pressure to reach 1 atm

• Boiling point elevation, ∆Tb = Kbcm

Kb = bp elevation constant, depends only on solvent

cm = molal concentration of solution

Freezing point depression

• Nonvolatile solutes will lower the freezing point of a solvent in a similar way to bp elevation

• ∆Tf = Kfcm

• Antifreeze both lowers the freezing point and raises the boiling point of the coolant

• Molecular weight of a solute can be determined by measuring its freezing point depression

Osmosis

• Semipermeable membrane: allows solvent molecules to pass but large solute molecules cannot

• Osmosis: flow of solvent through a semipermeable membrane to equalize solute concentrations on both sides of the membrane

• π = MRT (M = molar conc., R = gas constant, T = absolute temperature)

• Reverse osmosis: apply greater pressure to more concentrated solution and force pure solvent through membrane

Colligative properties of ionic solutions

• Effective concentration of ionic solutions is greater than molecular solutions even at the same molarity or molality

– Ionic compounds dissociate into individual ions• i = number of ions resulting from solvation of one formula unit• Multiply i in any colligative formula if the solute is ionic

∆Tf = iKfcm

∆P = iP°AXB

∆Tb = iKbcm

π = iMRT• i is only accurate in dilute solutions

Colloid formation

• Colloid: dispersion of particles throughout another substance or solution

– Differs from a solution in that its dispersed particles are more than one molecule in size (but still too small to see with the naked eye)

• Tyndall effect: colloids scatter light, while solutions do not

Types of colloids

• Aerosol: liquid or solid particles dispersed throughout a gas– Fog, smoke

• Emulsion: liquid droplets dispersed throughout another liquid– Particles of butterfat dispersed through homogenized milk

• Sol: solid particles dispersed in a liquid• Hydrophilic colloid: when there is a strong attraction between

particles and water– Gelatin

• Hydrophobic colloid: no attraction between particles and water

Coagulation and association

• Coagulation: particles of a colloid are made to aggregate and separate from solvent

– Milk curdles when its colloidal particles no longer have the same charge (they’re no longer repelled from each other)

• Association colloid: formed when colloidal particles have both hydrophobic and hydrophilic portions

– Soaps have long hydrocarbon chain (hydrophobic) and charged functionality (hydrophilic)

– The hydrophobic portions gather inwards to form spheres with a hydrophilic outside (micelles)