Chapter 12: Solutions

1. Chapter 12: Solutions Chemistry 1062: Principles of Chemistry II Andy Aspaas, Instructor

2. Solutions • Solution: homogeneous mixture of two or more substances (atoms, molecules, or ions) • Can exist in any state of matter

3. 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.

4. 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

5. 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

6. 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

7. 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)

8. 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

9. 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

10. 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

11. Henry’s law • Henry’s law: solubility of a gas is directly proportional to the partial pressure of the gas above the solution S = kHP where 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

12. 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)

13. 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

14. 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

15. 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

16. 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

17. 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

18. 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

19. 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

20. 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

21. 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)