Unit 2: Liquids & solids, solubility, equilibrium

1. Unit 2: Liquids & solids, solubility, equilibrium By: Ali Montgomery and Sam Block http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9018&vid=10000

2. General Properties of Aqueous Solutions • A substance such as NaCl whose aqueous solutions contain ions is an electrolyte • A substance such as C12H22O11 that does not form ions in solution is a nonelectrolyte. Most molecular substances are nonelectrolytes • Strong electrolytes are solutes that exist in solution almost entirely as ions • Weak electrolytes are solutes that exist in solution mostly as molecules with only a small fraction in the form of ions

3. Precipitation Reactions • Reactions resulting in the formation of an insoluble product are called precipitation reactions • A precipitate is an insoluble solid formed by a reaction in solution • The solubility of a substance is the amount of substance that can be dissolved in a given quantity of solvent at a specific temperature • Any substance with a solubility of less than 0.01 mol/L is generally considered insoluble http://www.iun.edu/~cpanhd/C101webnotes/chemical reactions/images/agcl.jpg

4. Solubility Rules (Appendix 2)http://dist113.org/dhs/Depts/Science/Hinton/appendix/a02.pdf • Rule 1ALKALI METALS and • AMMONIUM • Most alkali metal and ammonium salts are soluble. • All common salts of hydrogen, sodium, potassium, • and ammonium are soluble. • Rule 2ACETATES All acetates are soluble. Silver acetate is only • moderately soluble. • Rule 3CARBONATES Carbonates are insoluble except those covered by • Rule 1. All are soluble in the presence of acid. • Rule 4CHLORIDES, • BROMIDES, and • IODIDES • Most metal chlorides, bromides and iodides are • soluble except those of silver and lead. Lead (II) • chloride is moderately soluble.

5. Solubility Rules (continued) • Rule 5CHROMATES Chromates are insoluble except for those covered • by Rule 1 and those of manganese (II) and iron (III). • Chromates of mercury (II), calcium, and strontium • are slightly soluble. • Rule 6HYDROXIDES Hydroxides are insoluble except those covered by • Rule 1 and those of calcium, strontium, and barium. • Magnesium hydroxide is slightly soluble. • Rule 7NITRATES All nitrates are soluble. • Rule 8OXIDES Oxides are insoluble except those covered by Rule 1 • and those of calcium, strontium, and barium. • Oxides that are soluble usually react with water. • Rule 9PERMANGANATES Most permanganates are soluble. • Rule 10PHOSPHATES Phosphates are insoluble except those covered by • Rule 1. • Rule 11SULFATES Sulfates are soluble except those of strontium, • barium, and lead. Calcium and silver sulfates are • only moderately soluble. • Rule 12SULFIDES Sulfides are insoluble except those covered by • Rule 1.

6. Ionic Equations • Molecular Equations show the chemical formulas of reactants and products without indicating their ionic character • Ex. 3AgNO3(aq) + AlCl3(aq) 3AgCl(s) + Al(NO3)3(aq) • Complete ionic equations show all souble strong electrolytes as ions • Ex. 3Ag+ (aq) + 3NO3-(aq) + Al^3+(aq) + 3Cl-(aq) 3AgCl(s) + Al^3+(aq) + 3NO3-(aq) • Net ionic equations do not include spectator ions, which are quantities that can be canceled out • Ex. 3Ag+ (aq) + 3Cl-(aq) 3AgCl(s)

7. Equilibrium • Chemical Equilibrium occurs when opposing reactions proceed at equal rates • Although concentrations do not change at equilibrium, the reaction still occurs • For any equilibrium equation aA + bB <--> dD + eE, the equilibrium can be expressed by an equilibrium-constant expression as • The equilibrium constant expression depends only upon the stoichiometry of the reaction, not on its mechanism • Solids and liquids are only 1, and do not go into the constant because their concentrations don’t change • For an equilibrium constant expression in terms of pressure, Kp, use Kp=Kc(RT)^(Δn) to convert from the Kc constant • Δn is the moles of product-moles of reactant

8. More with Equilibrium Constants • If K>1, equilibrium lies to the right, and products are favored • If K<1, equilibrium lies to the left, and reactants are favored • The equilibrium-constant expression and the equilibrium constant of the reverse of a reaction are the reciprocals of those of the forward reaction • When a reaction is the sum of two or more reactions, its equilibrium constant is the product of the equilibrium constants for the individual reactions • Equilibrium constants vary with temperature.

9. InitialChangeEquilibrium! • If the equilibrium concentration of at least one species involved in a chemical reaction is known, we can use stoichiometry to calculate the other concentrations with an ICE chart. The steps to do this are: • 1) Put all known initial and equilibrium concentrations into an ice chart • 2) For species where both initial and equilibrium concentrations are provided, calculate the change in concentration • 3) Use stoichiometry of the reaction (keeping in mind the effects of coefficients on the species involved) to calculate the change in concentration for the other species involved • 4) Use the change in concentration to arrive at the equilibrium concentrations. http://images.google.com/imgres?imgurl=http://dclips.fundraw.com/zobo500dir/ice_cube_jarno_vasamaa_.jpg&imgrefurl=http://www.fundraw.com/clipart/clip-art/00001488/Melting-Ice-Cube/&usg=__bkfO82OyKEtyWi3D45nS

10. Reaction Quotients • If the Reaction Quotient Q (a number obtained by substituting actual product and reactant concentrations into an equilibrium-constant equation) equals K, the system is at equilibrium • If Q>K, the concentration of products is too big so the reaction shifts left to offset it by increasing the concentration of reactants • If Q<K, the concentration of reactants is too big, so the reaction shifts right, which increases the amount of products

11. Le Chatelier’s Principle • Changes in temperature, pressure (or volume), or the concentration of one substance involved in the reaction shift the equilibrium position to offset the changes; catalysts do not shift equilibrium http://myphlip.pearsoncmg.com/phproducts/student/ http://wps.prenhall.com/wps/media/access/PearsonDefault/3064/3137997/login.html

12. Le Chatelier’s Principle • Adding a substance causes the reaction to shift as more of that excess substance is consumed. Removing a substance causes the reaction to shift to produce more that substance • Reducing the volume (increasing the pressure) of an equilibrium with gaseous components causes the system to shift to the side that reduces the number of moles; decreasing the pressure (increasing volume) shifts the system to the side that increases the number of moles of gas • If the system is exothermic, heat is treated as a product. Raising the temperature shifts the system left • If the system is endothermic, heat is a reactant. Raising the temperature shifts the system right

13. Solubility Product • The solubility product, Ksp, equals the product of the concentration of the ions involved in the equilibrium, raised to the power of the coefficient in the equation • BaSO4 (s) <-> Ba ^2+ (aq) + SO4 ^2- (aq) • Ksp=[Ba^2+][SO4 ^2-] • The Ksp can be used to calculate the solubility of a compound

14. Factors that Affect Solubility • The presence of common ions reduces solubilty • Temperature affects solubility of substances dissolved in water • The solubility of compounds containing basic anions increases as the solution is made more acidic, and pH decreases • Salts with anions of negligible basicity, anions of strong acids such as Cl-, are not affected by pH changes

15. Precipitation • If Q>Ksp, precipitation occurs until Q=Ksp • If Q=Ksp, equilibrium exists in a saturated solution • If Q<Ksp, solid dissolves until Q=Ksp

16. Intermolecular Forces • Gas molecules undergo constant, chaotic motion • Liquid molecules are free to move, but kept in close proximity • Solid molecules have strong enough attractive forces to restrain molecular motion

17. IMFA’s Continued • Dipole-dipole forces • Dispersion force depends on polarizability, size, and shape • London dispersion forces • All molecules • Especially when molecules very close together • Hydrogen bonding • exist between a hydrogen atom in a polar bond and an electronegative element, (H-O, H-F, H-N)

18. Flow • Viscosity: Resistance to flow! • Greater weight, attraction increased viscosity • Surface tension: Energy required to increase the surface area of a liquid by a unit amount • Adhesive: Bind substance to surface • Cohesive: Bind similar molecules together • Capillary action: Rise of liquids up narrow tube

19. (vaporization) Phase Changes In J, must convert to kJ http://wps.prenhall.com/esm_brown_chemistry_10_upgrade/47/12261/3138864.cw/index.html

20. Phase Diagram http://wps.prenhall.com/esm_brown_chemistry_10_upgrade/47/12261/3138864.cw/index.html

21. Crystalline Structures • Simple cubic: 1 atom, V=8r3, e=2r • Body-centered: 2 atoms, V=(4r/√3)3, e=4r/√3 • Face-centered: 4 atoms, V= (32r3/√2), e=4r/√2 http://www.substech.com/dokuwiki/lib/exe/fetch.php?w=&h=&cache=cache&media=crystal_latti ce.png

22. James Bonding in Solids • Molecular solids: atoms/molecules held together by intermolecular forces (London dispersion, dipole-dipole, hydrogen bonds), soft, relatively low boiling points, poor thermal and electrical conduction, ex: methane, sucrose, and dry ice • Covalent-network solids: atoms held together by covalent bonds, very hard, high melting points, poor thermal and electrical conductors, ex: diamonds • Ionic solids: ions held together by ionic bonds, hard and brittle, high melting points, poor thermal and electrical conduction, ex: salts • Metallic solids: metal atoms held together by metallic bonds, vary in strength of bonding, wide range of physical properties (hardness, melting points), malleable and ductile, excellent thermal and electrical conductors, ex: copper, iron, aluminum

23. Saturated Solutions and Solubility • Saturated: Solution that is in equilibrium with undissolved solute • Solubilty: Amount of solute needed to form a saturated solution in a given amount of solvent • Unsaturated: Less solute than is needed to form a saturated solution • Superaturated: Not unsaturated. The opposite, in fact

24. Henry’s Law (It’s French) • Sg=kPg • Sg is the solubility of the gas in the solution phase • Pg is the partial pressure of the gas over the solution • k is a proportionality constant

25. Concentration Stuff • Mass %=Mass of component is solution/ total mass solution • PPM: Parts per million (10^6) • Molarity=Moles solute / liters solution • Molality= moles of solute / kg solvent

26. Colligative Properties • Boiling-point elevation: • Freezing-point depression: • Osmosis: pi=(n/v)RT=MRT • The net movement of solvent is always toward the solution with the higher solute concentration

27. Raoult’s Law • Pvapor=XP°vapor • X is the mole fraction of a solvent in solution • P°vapor is the vapor pressure of the pure solvent • Pvapor is the partial pressure of a solvent over a solution • = is the equal sign

28. Vocab • Volatile: Liquids that evaporate readily • Miscible: Pairs of liquid that mix in all proportions • Journalism: No science involved!