1 / 15

Unit 2: Liquids and solids, solubility, equilibrium

Unit 2: Liquids and solids, solubility, equilibrium. Will Barkalow and Price Ryan. Intermolecular forces. In order from weakest to strongest London Dispersion – non-polar/non-polar Dipole-Dipole – polar/polar Hydrogen Bond – ultra polar/ultra polar H with N,O,F metalli c ionic

hester
Download Presentation

Unit 2: Liquids and 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 and solids, solubility, equilibrium Will Barkalow and Price Ryan

  2. Intermolecular forces • In order from weakest to strongest • London Dispersion – non-polar/non-polar • Dipole-Dipole – polar/polar • Hydrogen Bond – ultra polar/ultra polar • H with N,O,F • metallic • ionic • Covalent Bond (network) – certain C-family elements and compounds

  3. Cubic crystal structure • Simple Cubic (1 atom) • V = e3 = 8r3 • Body-Centered Cubic (2 atoms) • 4r = e • V = e3 =(4r/ )3 • Face-Centered Cubic (4 atoms) • V = e3 = (4r/ )3

  4. Phase Changes • Matter exists as solids, liquids, or gases • Matter changes between these three phases based on temperature or atmospheric pressure

  5. Phase diagram • Phase Diagrams – show which state matter exist at certain temp and pressure • Triple Point (T) – all three phases are at equilibrium • Critical Point (C) – highest temp of pressure where a distinct gas and liquid phase can exist • Supercritical Fluid (SCF) – liquid and gas phases are indistinguishable

  6. Solubility • The amount of substance that can be dissolved in a solvent at a given temperature (g/L) • Some Simple Solubility Rules • NO3- is always soluble • C2H3O2- is always soluble • OH- is insoluble with everything except alkali metals, NH4+, Ca2+, Sr2+, and Ba2+ • PO43- is insoluble with everything except alkali metals and NH4+

  7. Equilibrium • aA + bBcC + dD • Occurs when opposing reactions proceed at equal rates with constant concentrations • KC = equilibrium constant • KC = {[C]c[D]d}/{[A]a[B]b} • No Solids and Liquids included in equation • Q = Reaction Quotient • Q = {[C]c[D]d}/{[A]a[B]b} • When not at equilibrium • Q < K  Shift Right, Q > K  Shift Left, Q = K  Equilibrium Established • No solids or liquids included in equation

  8. Equilibrium continued • KP = KC(RT)-n • or same expression as Kc, but using pressures • KSP= solubility product constant, this indicates how soluble a solid is in water • KSP = (C)c(D)d • Reactants are Solids so they are not included in the equation

  9. LeChâtlier’s Principle • A system in equilibrium, when disturbed, will shift to the extent necessary to restore equilibrium • Example: • N2 + 3H2 2NH3 + 94K • Disturbance add N2: reactant is increased so product is increased, shift right • Disturbance increase temp: tries to lower energy, lower products, shift left • Disturbance increase pressure: will favor side with fewer moles, fewer moles on product side, increase products, shift right • Disturbance add catalyst: reaction goes faster but no change to each side, no shift

  10. Predicting Formation of precipitates based on Ksp • Knet= Ksp x Kf • Kf = formation constant • Used in reactions with complex ions • Lower Ksp means it’s less soluble. The combination with a complex ion makes it more soluble (the Kf value times the Kspvalue is much larger than the original Kspvalue).

  11. Concentration Units • Cmolar = nsolute/Vsolution • Cmolal = nsolute /msolvent (kg) • Mass Percent = msolute /msolution • Mole Fraction = nsolute /nsolution • Volume Fraction = Vsolute /Vsolution • Parts Per Million (PPM) = msolute (mg) / msolvent (kg) ≈ msolute (mg) / msolution (kg) when mass of solute is small enough

  12. Henry’s Law • Sg = kHPg • The solubility of gas increases in direct proportion to the partial pressure above the solution

  13. ColligativeProperties • Boiling Point Elevation: Tb = iCmolalKb • i = Van’t Hoff Factor – when dissolved , # of particles the molecule breaks into (Ex. NaCl = 2) • Kb = Boiling point elevation constant • solutes make it harder for water molecules to vaporize, thus requiring more energy to boil • Freezing Point Depression: Tf = iCmolalKf • Kf = freezing point depression constant • solutes make it harder for water molecules to form an orderly crystalline structure

  14. Raoult’s Law (Vapor Pressure) • PA = XAPOA • PA = solution vapor pressure • XA = mole fraction of solvent (nsolvent/nsolution) • POA = solvent vapor pressure • Adding a non-volatile solute to a solution will lower the vapor pressure, thus raising the boiling point

  15. Osmosis • V = nRT •  = nRT/V = CmolarRT •  = osmotic pressure • R = 8.314 if pressure is in Pascals (Pa). R = 0.08206 if pressure is in atm • Pressure required to prevent osmosis by a solvent toward a solution with a higher solute concentration

More Related