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Overview of Ch 11-13

Overview of Ch 11-13. Properties of Solutions Chapter 11. Solution Composition. 1. Molarity ( M ) = 2. Mole fraction (  A ) = 3. Molality ( m ) =. Henry’s Law. The amount of a gas dissolved in a solution is directly proportional to the pressure of the gas above the solution. P = k C

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Overview of Ch 11-13

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  1. Overview of Ch 11-13

  2. Properties of SolutionsChapter 11

  3. Solution Composition 1. Molarity (M) = 2. Mole fraction (A) = 3. Molality (m) =

  4. Henry’s Law The amount of a gas dissolved in a solution is directly proportional to the pressure of the gas above the solution. P = kC P = partial pressure of gaseous solute above the solution C = concentration of dissolved gas k = the Henry’s Law constant

  5. Temperature Effects • Solubility of gases generally decreases with temperature. • Solubility of solids generally increases with temperature.

  6. Colligative Properties Depend only on the number, not on the identity, of the solute particles in an ideal solution. • Vapor pressure depression • Boiling point elevation • Freezing point depression • Osmotic pressure increase

  7. Raoult’s Law The presence of a nonvolatile solute lowers the vapor pressure of a solvent. Psoln= solventPsolvent Psoln= vapor pressure of the solution solvent= mole fraction of the solvent Psolvent= vapor pressure of the pure solvent

  8. Boiling Point Elevation A nonvolatile solute elevates the boiling point of the solvent. T = Kbmsolute Kb = molal boiling point elevation constant m = molality of the solute

  9. Freezing Point Depression A nonvolatile solute depresses the freezing point of the solvent. T = Kfmsolute Kf= molal freezing point depression constant m = molality of the solute

  10. Osmotic Pressure Osmosis: The flow of solvent into the solution through a semipermeable membrane. Osmotic Pressure: A nonvolatile solute increases the osmotic pressure of the solvent.

  11. Chemical EquilibriumChapter 13 • The state where the concentrations of all reactants and products remain constant with time.

  12. Equilibrium Constant jA + kB lC + mD • The equilibrium expression:

  13. 4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(g)

  14. Manipulations of K • The equilibrium constant for a reaction is the reciprocal of that for the reaction written in reverse. • When the equation for a reaction is multiplied by n, Knew = (Koriginal)n

  15. K v. Kp • For jA + kB lC + mD Kp = K(RT)n • n = sum of coefficients of gaseous products minus sum of coefficients of gaseous reactants.

  16. Heterogeneous Equilibria • . . . are equilibria that involve more than one phase. CaCO3(s)  CaO(s) + CO2(g) K = [CO2] • The position of a heterogeneous equilibrium does not depend on the amounts of pure solids or liquids present.

  17. Reaction Quotient • . . . helps to determine the direction of the move toward equilibrium. • The law of mass action is applied with initial concentrations.

  18. H2(g) + F2(g)  2HF(g) • Q < K, shift right • Q > K, shift left

  19. Solving Equilibrium Problems 1. Write the equilibrium expression. 2. Set up an “ICE” box with relevant concentrations. 3. Use the stoichiometry of the reaction to determine changes in products and reactants, solving for unknowns.

  20. Le Châtelier’s Principle • . . . if a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction that tends to reduce that change.

  21. Effects of Changes on the System 1.Concentration: The system will shift away from the added component. 2. Temperature: treat the energy change as a reactant (endothermic) or product exothermic).

  22. Effects of Changes on the System (continued) 3. Pressure: a. Addition of inert gas does not affectthe equilibrium position. b.Decreasing the volume shifts the equilibrium toward the side with fewer moles.

  23. Chemical KineticsChapter 12 • The area of chemistry that concerns reaction rates.

  24. Reaction Rate • Change in concentration(conc) of a reactant or product per unit time. Reaction rates are positive by convention.

  25. (Differential) Rate Laws Rate = k[NO2]n • k = rate constant • n = rate order

  26. Types of Rate Laws • Differential Rate Law: expresses how rate depends on concentration. • Integrated Rate Law: expresses how concentration depends on time.

  27. Method of Initial Rates • Initial Rate: the “instantaneous rate” just after the reaction begins. • The initial rate is determined in several experiments using different initial concentrations.

  28. Overall Reaction Order • Sum of the order of each component in the rate law. rate = k[H2SeO3][H+]2[I]3 • The overall reaction order is 1 + 2 + 3 = 6.

  29. First-Order Rate Law For aA  Products in a 1st-order reaction, • Integrated first-order rate law is ln[A] = kt + ln[A]o

  30. Half-Life of a 1st-Order Rxn • t1/2= half-life of the reaction • k = rate constant • For a first-order reaction, the half-life does not depend on concentration.

  31. Second-Order Rate Law • For aA  products in a second-order reaction, • Integrated rate law is:

  32. Half-Life of a 2nd-Order Rxn • t1/2= half-life of the reaction • k = rate constant • Ao= initial concentration of A • The half-life is dependent upon the initial concentration.

  33. Zero-Order Rate Law • For aA  products in a zero-order reaction, Rate= k • Integrated rate law is [A] = -kt + [A]o

  34. Half-Life of a Zero-Order Rxn t1/2 = [A]o • t1/2= half-life of the reaction • k = rate constant • [A]o= initial concentration of A • The half-life is dependent upon the initial concentration. 2k

  35. Reaction Mechanism • The series of stepsby which a reaction occurs. • A chemical equation does not tell us how reactants become products - it is a summary of the overall process.

  36. Reaction Mechanism (continued) • The reaction has many steps in the reaction mechanism.

  37. Intermediate: formed in one step and used up in a subsequent step and so is never seen as a product. • Molecularity: the number of species that must collide to produce the reaction indicated by that step. • Elementary Step: A reaction for which a rate law can be written from its molecularity. • uni, bi and termolecular

  38. Rate-Determining Step • In a multistep reaction, it is the slowest step. It therefore determines the rate of reaction.

  39. Arrhenius Equation • Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy). • Orientation of reactants must allow formation of new bonds.

  40. k = rate constant • A = frequency factor • Ea = activation energy • T= temperature (in K) • R= gas constant

  41. lnk= -Ea 1 + ln A R T slope = -Ea /R • catalysts decrease Ea.

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