1 / 28

Unit 2: Liquids & solids, solubility, equilibrium

Unit 2: Liquids & solids, solubility, equilibrium. By: Ali Montgomery and Sam Block. http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9018&vid=10000. General Properties of Aqueous Solutions. A substance such as NaCl whose aqueous solutions contain ions is an electrolyte

hea
Télécharger la présentation

Unit 2: Liquids & solids, solubility, equilibrium

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  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!

More Related