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First, system reaches equilibrium

Chemical Equilibrium. Introduction 1.) Equilibria govern diverse phenomena Protein folding, acid rain action on minerals to aqueous reactions 2.) Chemical equilibrium applies to reactions that can occur in both directions: reactants are constantly forming products and vice-versa

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First, system reaches equilibrium

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  1. Chemical Equilibrium • Introduction • 1.)Equilibria govern diverse phenomena • Protein folding, acid rain action on minerals to aqueous reactions 2.) Chemical equilibrium applies to reactions that can occur in both directions: • reactants are constantly forming products and vice-versa • At the beginning of the reaction, the rate that the reactants are changing into the products is higher than the rate that the products are changing into the reactants. • When the net change of the products and reactants is zero the reaction has reached equilibrium. Then, system continually exchanges products and reactants, while maintaining equilibrium distribution. First, system reaches equilibrium Product Reactants At equilibrium the amount of reactants and products are constant, but not necessarily equal

  2. Chemical Equilibrium • Equilibrium Constant • 1.)The relative concentration of products and reactants at equilibrium is a constant. 2.) Equilibrium constant (K): • For a general chemical reaction Equilibrium constant: Where: - small superscript letters are the stoichiometry coefficients - [A] concentration chemical species A relative to standard state

  3. Chemical Equilibrium • Equilibrium Constant 2.) Equilibrium constant (K): • A reaction is favored when K > 1 • K has no units, dimensionless - Concentration of solutes should be expressed as moles per liter (M). - Concentrations of gases should be expressed in bars. ►express gas as Pgas, emphasize pressure instead of concentration ►1 bar = 105 Pa; 1 atm = 1.01325 bar - Concentrations of pure solids, pure liquids and solvents are omitted ►are unity ►standard state is the pure liquid or solid 3.)Manipulating Equilibrium Constants Consider the following reaction: Reversing the reaction results in a reciprocal equilibrium reaction:

  4. Chemical Equilibrium Equilibrium Constant 3.)Manipulating Equilibrium Constants If two reactions are added, the new K is the product of the two individual K values: K1 K2 K3

  5. Chemical Equilibrium • Equilibrium Constant 3.)Manipulating Equilibrium Constants • Example: Given the reactions and equilibrium constants: Kw= 1.0 x 10-14 KNH3= 1.8 x 10-5 Find the equilibrium constant for the reaction: Solution: K1= Kw K2=1/KNH3 K3=Kw*1/KNH3=5.6x10-10

  6. Chemical Equilibrium • Le Châtelier’s Principal 1.) What Happens When a System at Equilibrium is Perturbed? • Change concentration, temperature, pressure or add other chemicals • Equilibrium is re-established • Reaction accommodates the change in products, reactants, temperature, pressure, etc. • Rates of forward and reverse reactions re-equilibrate

  7. Chemical Equilibrium • Le Châtelier’s Principal 1.) What Happens When a System at Equilibrium is Perturbed? • Le Châtelier’s Principal: - the direction in which the system proceeds back to equilibrium is such that the change is partially offset. Consider this reaction: At equilibrium: To return to equilibrium (balance), some (not all) CO and H2 are converted to CH3OH Add excess CO(g): If all added CO was converted to CH3OH, then reaction would be unbalanced by the amount of product

  8. Chemical Equilibrium • Le Châtelier’s Principal 2.) Example: Consider this reaction: At one equilibrium state:

  9. Chemical Equilibrium • Le Châtelier’s Principal 2.) Example: What happens when: According to Le Châtelier’s Principal, reaction should go back to left to off-set dichormate on right: Use reaction quotient (Q), Same form of equilibrium equation, but not at equilibrium:

  10. Chemical Equilibrium • Le Châtelier’s Principal 2.) Example: Because Q > K, the reaction must go to the left to decrease numerator and increase denominator. Continues until Q = K: 1. If the reaction is at equilibrium and products are added (or reactants removed), the reaction goes to the left 2. If the reaction is at equilibrium and reactants are added ( or products removed), the reaction goes to the right

  11. Chemical Equilibrium • Le Châtelier’s Principal 3.) Affect of Temperature on Equilibrium 1. Equilibrium constant of an endothermic reaction (DHo = +) increases if the temperature is raised. 2. Equilibrium constant of an exothermic reaction (DHo = -)decreases if the temperature is raised. D DH = + D DH = -

  12. Chemical Equilibrium • Solubility Product 1.) Equilibrium constant for the reaction which a solid salt dissolves to give its constituent ions in solution • Solid omitted from equilibrium constant because it is in a standard state • Example:

  13. Chemical Equilibrium • Solubility Product 1.) Saturated Solution – contains excess, undissolved solid • Solution contains all the solid capable of dissolving under the current conditions • EXAMPLE: Find [Cu2+] in a solution saturated with Cu4(OH)6(SO4) if [OH-] is fixed at 1.0x10-6M. Note that Cu4(OH)6(SO4) gives 1 mol of SO42- for 4 mol of Cu2+?

  14. Chemical Equilibrium • Solubility Product 2.) If an aqueous solution is left in contact with excess solid, the solid will dissolve until the condition of Ksp is satisfied • Amount of undissolved solid remains constant • Excess solid is required to guarantee ion concentration is consistent with Ksp 3.)If ions are mixed together such that the concentrations exceed Ksp, the solid will precipitate. 4.) Solubility product only describes part of the solubility of a salt • Only includes dissociated ions • Ignores solubility of solid salt

  15. Chemical Equilibrium Common ion effect – a salt will be less soluble if one of its constituent ions is already present in the solution. Decrease in the solubility of MgF2 by the addition of NaF PbCl2 precipitate because the ion product is greater than Ksp.

  16. Chemical Equilibrium • Common Ion Effect 1.) Affect of Adding a Second Source of an Ion on Salt Solubility • Equilibrium re-obtained following Le Châtelier’s Principal • Reaction moves away from the added ion • EXAMPLE: Find [Cu2+] in a solution saturated with Cu4(OH)6(SO4) if [OH-] is fixed at 1.0x10-6M and 0.10M Na2SO4 is added to the solution.

  17. Chemical Equilibrium • Complex Formation 1.) High concentration of an ion may redissolve a solid • Ion first causes precipitation • Forms complex ions, consists of two or more simple ions bonded to each other ppt. formation Complex forms and redissolves solid

  18. Chemical Equilibrium • Complex Formation 2.) Lewis Acids and Bases • M+ acts as a Lewis acid  accepts a pair of electrons • X- acts as a Lewis base  donates a pair of electrons • Bond is a coordinate covalent bond adduct ligand Lewis base Lewis acid

  19. Chemical Equilibrium • Complex Formation 3.) Affect on Solubility • Formation of adducts increase solubility • Solubility equation becomes a complex mixture of reactions - don’t need to use all equations to determine the concentration of any species Ksp Implies low Pb2+ solubility: Only one concentration of Pb2+ in solution Concentration of Pb2+ that satisfies any one of the equilibria must satisfy all of the equilibria All equilibrium conditions are satisfied simultaneously

  20. Chemical Equilibrium • Complex Formation 3.) Affect on Solubility • Total concentration is dependent on each individual complex species Total solubility of lead depends on [I-] and the solubility of each individual complex formation.

  21. Chemical Equilibrium • Complex Formation 3.) Affect on Solubility • EXAMPLE: Given the following equilibria, calculate the concentration of each zinc-containing species in a solution saturated with Zn(OH)2(s) and containing [OH-] at a fixed concentration of 3.2x10-7M. Zn(OH)2 (s) Ksp = 3.0x10-16 Zn(OH)+b1 = 2.5 x104 Zn(OH)3-b3 = 7.2x1015 Zn(OH)42-b4 = 2.8x1015

  22. Chemical Equilibrium • Acids and Bases 1.)Protic Acids and Bases – transfer of H+ (proton) from one molecule to another • Hydronium ion (H3O+) – combination of H+ with water (H2O) • Acid – is a substance that increases the concentration of H3O+ • Base – is a substance that decreases the concentration of H3O+ - base also causes an increase in the concentration of OH- in aqueous solutions 2.) Brønsted-Lowry – definition does not require the formation of H3O+ • Extended to non-aqueous solutions or gas phase • Acid – proton donor • Base – proton acceptor acid acid base salt

  23. Chemical Equilibrium • Acids and Bases 3.)Salts – product of an acid-base reaction • Any ionic solid • Acid and base neutralize each other and form a salt • Most salts with a single positive and negative charge dissociate completely into ions in water 4.) Conjugate Acids and Bases Products of acid-base reaction are also acids and bases A conjugate acid and its base or a conjugate base and its acid in an aqueous system are related to each other by the gain or loss of H+

  24. Chemical Equilibrium • Acids and Bases 5.)Autoprotolysis – acts as both an acid and base • Extent of these reactions are very small water • - H3O+ is the conjugate acid of water • - OH- is the conjugate base of water • Kw is the equilibrium constant for the dissociation of water Acetic acid

  25. Chemical Equilibrium • Acids and Bases 6.)pH – negative logarithm of H+ concentration • Ignores distinction between concentration and activities (discussed later) • A solution is acidic if [H+] > [OH-] • A solution is basic if [H+] < [OH-] • An aqueous solution has a neutral pH if [H+]=[OH-] - This occurs when [H+] = [OH-] = 10-7M or pH = 7

  26. Chemical Equilibrium • Acids and Bases 6.)pH • Example: What is the pH of a solution containing 1x10-6 M H+? What is [OH-] of a solution containing 1x10-6 M H+?

  27. Chemical Equilibrium • Acids and Bases 7.) Strengths of Acids and Bases • Depends on whether the compound react nearly completely or partially to produce H+ or OH- • strong acid or base completely dissociate in aqueous solution - equilibrium constants are large - everything else termed weak Strong  no undissociated HCl or KOH

  28. Chemical Equilibrium • Acids and Bases 7.) Strengths of Acids and Bases • weak acids react with water by donating a proton - only partially dissociated in water - equilibrium constants are called Ka – acid dissociation constant - Ka is small • weak bases react with water by removing a proton - only partially dissociated in water - equilibrium constants are called Kb – base dissociation constant - Kb is small Ka Equivalent Ka Kb Equivalent Kb

  29. Chemical Equilibrium Some Common Weak Acids (carboxylic acids)

  30. Chemical Equilibrium Some Common Weak Acids (Metals cations)

  31. Chemical Equilibrium Some Common Weak Bases (amines) • The Ka or Kb of an acid or base may also be written in terms of “pKa” or “pKb” • As Ka or Kb increase  pKa or pKb decrease - a strong acid/base has a high Ka or Kb and a low pKa or pkb

  32. Chemical Equilibrium • Acids and Bases 8.) Polyprotic Acids and Bases – can donate or accept more than one proton • Ka or Kb are sequentially numbered - Ka1,Ka2,Ka3 Kb1,Kb2,Kb3

  33. Chemical Equilibrium Acids and Bases 8.) Relationship Between Ka and Kb

  34. Chemical Equilibrium • Acids and Bases 8.) Relationship Between Ka and Kb • EXAMPLE: Write the Kb reaction of CN-. Given that the Ka value for HCN is 6.2x10-10, calculate Kb for CN-.

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