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Chemistry chapter 20


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Chemistry chapter 20

  1. 1. By Logan Danielson
  2. 2. TitrationsTitrations are done to determine the unknown concentrations of unknown acids.It is based of off determining how much known acid or base is needed to reach the equivalence point of the unknown liquid and at what pH the equivalence point exists.The equivalence point is the middle point of the fastest changing pH area. Usually a color indicator is used to help find the equivalence point as well as a pH indicator.
  3. 3. Solubility EquilibriaAdding an ionic solid to water It will be in equilibrium so long as some solid exists in solutionMg(OH)2(s)<-> Mg2+(aq) + 2 OH-(aq) Ksp= [Mg2+][OH-]2  This is the solubility product constant for the above equation  Note, the solid is not part of the equation
  4. 4. Predicting precipitationFor Mg(OH)2(s)<-> Mg2+(aq) + 2 OH-(aq) If Q= [Mg2+][OH-]2 < Ksp then the there will be no precipitate. If Q > Ksp then the there will be a precipitate. Note, if Q was < Ksp then only Mg2+(aq) or OH-(aq) would need to be added to cause precipitation.
  5. 5. ThermodynamicsIs the flow of energy1st law of thermodynamics  The total energy of the universe is constant, energy can not be created nor destroyed  If you decrease the energy in the system, then you increase the energy of the surrounding  If you increase the energy in the system, then you decrease the energy of the surrounding U = internal energy = all potential energy + all kinetic energy dU universe = dU system + dU surroundings  dU system = - dU surroundings
  6. 6. Thermodynamics 2: heat and workdU system = Heat (q) + Work (w)  q= Heat = energy flow from the change in temperature  w= Work = energy flow from movement against a forceClosed system- do not allow matter in or out of the systemq<0 – heat flows out of the systemq>0 – heat flows into the systemw<0 – the system does work on the surroundingw>0 – the surrounding does work on the system
  7. 7. Thermodynamics 3: enthalpyThe most common work in chemical systems w= -P dV  Work= - Pressure * change in Volume dU system = q – P * dV dH (enthalpy) = dU + d(PV) dH = dU + P*dH @ constant P Plugging in “dU system = q – P * dV” you get: dH = q - P*dH + P*dH = q @ constant P dH = q @ constant P
  8. 8. 2nd law of thermodynamicsAll spontaneous processes increase the entropy of the universe dS (entropy) > 0 for spontaneous processesEntropy is a measure of the dispersive-ness of energy A larger value of entropy is a more dispersed energy
  9. 9. Guessing thermodynamicsC6H12O6(s) -> 2 C2H3OH(s) + 2 CO2(g) dH>0 because we are breaking more bonds then we are forming dS>0 because there are more moles on the rightH2O(s) -> H2O(l) dS>0 because liquid is more dispersible than solidH2O(l) -> H2O(g) dS>0 because gas is more dispersible than liquid
  10. 10. Entropy guessing rulesIncrease Entropy Breaking bonds without making new ones Change to a favored phase More moles on the product side
  11. 11. EntropydS universe = (dS system + dS surroundings) > 0dS surroundings = -q / TdS system - q / T > 0 At constant P, q=dH so:  dS - dH / T > 0 systemFor phase changes: dS system - dH / T is approximately 0, so:  dS system = dH / T
  12. 12. Calculating entropyaA + bB -> cC + dD dS reaction = c*SCo + d*SDo – a*SAo – b*S Bo SAo is the standard entropy of A
  13. 13. Gibbs free energy G = free energy dG = dH – TdS @ constant P and T dG < 0 for a spontaneous process Reactions that are dG>0 won’t happen, but their reverse reaction will Reactions that are dG=0 will have nothing happen because the system is at equilibrium The free energy and the equilibrium constant are related:  dGo= - RT ln K  dGo is the dG at standard temperature and pressure  R is the gas law constant, R=8.314 J/(K*mol)  T is the temperature  ln K is the natural logarithm of the equilibrium constant
  14. 14. Gibbs free energydG = dH – T*dS @ constant P and TdHo dSo dGo>0, endothermic >0, increase in energy (+)-T(+), more negative at higher temperatures<0, exothermic <0, decrease in energy (-)-T(-), more negative at lower temperatures>0, endothermic <0, decrease in energy (+)-T(-), always positive, never spontaneous<0, exothermic >0, increase in energy (-)-T(+), always negative, always spontaneousAt low T: dG is approximately= dHAt high T: dG is approximately= – T*dS