1 / 11

Kinetics and Equilibria

Kinetics and Equilibria. By definition, kinetic processes are not equilibrium processes . In fact, we may think of kinetic processes as the mechanism that nature uses to reach the equlibrium state. Binary Collision rate in forward direction. Binary Collision rate in reverse direction.

ochoaj
Télécharger la présentation

Kinetics and Equilibria

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. Kinetics and Equilibria By definition, kinetic processes are not equilibrium processes. In fact, we may think of kinetic processes as the mechanism that nature uses to reach the equlibrium state. Binary Collision rate in forward direction. Binary Collision rate in reverse direction. has 2 rate constants, we can write, assuming these are ELEMENTARY reaction steps: (Equilibrium condition) Where [A]e etc. are the equilibrium concentrations of [A] etc.

  2. - E / RT - E / RT = k = A e k A e Ar Af r r f f Using the Arrhenius form for the rate constants kf and kr But later we will learn (or you already know from high school): G0=H0 - T S0 ln[Keq]= -G0/RT Where H0 is the enthalpy change for the reaction and S0 is the entropy change for the reaction. G0 is called the Free Energy

  3. Equating these two forms for the equilibrium constant allows us to connect thermodynamics and kinetics!  Kinetic form of Keq Thermodynamic form of Keq “Identify” Af / Ar with {e (S/R)} (T “independent” assuming ∆Sº indep of T). Act State EAf EAr A + B +∆Ho = EAf - EAr (∆Ho = Enthalpy change for A+B C+D) C + D

  4. Acid-Base Equilibria Several ways to define acid, base: 1) HCl  H+ + Cl- CH3COOH  CH3COO- + H+ NaOH  Na+ (aq) + OH- (aq) Oversimplified: H+ ~ 10-13 cm in diameter because is a free proton (unique in + charged species) Evidence exists for presence of H3O+ in solution. Small size of H+ allows it to be incorporated into the structure of the solvent.

  5. H3O+ is called the hydronium ion-is particularly stable. Less evidence for species like H9O4+ (4H2O+H+) More accurate view: Some non-OH- species can neutralize acids: HCl(aq)+NH3 NH4++Cl- 2) HCl (aq) + H2O  H3O+ (aq) + Cl- (aq) acid base acid base Note: H2O is a base here.

  6. CO32- + H2O  HCO3- + OH- Conjugate pairs: CO32-, HCO3- and H2O, OH- Strengths of Acids and Bases: Equate to tendency to transfer a proton to H2O HCl + H2O  H3O+ + Cl- HSO4- + H2O  H3O+ + SO42- Need a quantitative measure of acidity or H+ donating power.

  7. HA + H2O  H3O+ + A- Note that large K is associated with strong acid since it means numerator is large compared to denominator. Large KHA is a good proton donor to H2O. If acid is strong, e.g., HCl, then conjugate base is weak (Cl-) Will prove this latter in a quantitative fashion. 3) Lewis Concept (most general):Acid is any substance that can accept electrons and a base is any substance that can donate electrons.

  8. Acid ionization constant Acid-Base Equilibria Considerations (K’ is a true equilibrium constant) Ka is called the acid ionization constant NH3 + H2O  NH4+ + OH -

  9. pH= -log10[H3O+]

  10. Bonus * Bonus * Bonus * Bonus * Bonus * Bonus

  11. Suppose you have [H3O+] from added HCl = 0.1 M = 10-1 M Then [10 -1][OH-] = 10-14 In pure H2O, [OH-] = [H3O+] = 10-7, but Weak Acids and Bases Add acetic acid to H2O: K=Ka What are [CH3COOH], [H3O+], [CH3COO-]?

More Related