1 / 36

Chemical Equilibria

Chemical Equilibria. Professor Brian Kinsella. The Law of Mass Action. A + B ↔ C + D The velocity at which A and B react is proportional to their concentrations ν 1 = k 1 x [A] x [B] ν 2 = k 2 x [C] x [D]

guillermo
Download Presentation

Chemical 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. Chemical Equilibria Professor Brian Kinsella

  2. The Law of Mass Action • A + B ↔ C + D • The velocity at which A and B react is proportional to their concentrations • ν1 = k1 x [A] x [B] • ν2 = k2 x [C] x [D] • At equilibrium the velocities of the forward and reverse reaction will be equal and ν1 = ν2 • k1 x [A] x [B] = k2 x [C] x [D]

  3. The Law of Mass Action • Or • K = The equilibrium constant for the reaction at a given temperature • For a reversible reaction the equation may be generalised

  4. The Law of Mass Action (X) indicates the concentration of the reactants and products, but to be strictly correct it is the activity of reactants and products that should be used.

  5. Activity and Activity Coefficient • For a binary electrolyte • AB ↔ A+ + B- • activity = (concentration) x (activity coefficient) • Thus at any molar concentration

  6. Activity and Activity Coefficient • This is the rigorously correct expression for the law of mass action as applied to weak electrolytes. • The activity coefficient varies with concentration and ionic strength (IS). For ions it varies with the valency and is the same for all dilute solutions having the same ionic strength. • An increase in IS causes the activity coefficient and activity to decrease.

  7. Calculation of Ionic Strength • The ionic strength for 0.1 M HNO3 and 0.2M Ba(NO3)2 = 0.5{(0.1 x 12 + 0.1 x 12)HNO3 + (0.2 x 22 + 0.2 x 2 x 12)} = 0.5{0.2 + (0.8 + 0.4)} = 0.5{0.2 + 1.2} = 0.7 • The activity coefficient of unionised molecules do not differ considerably from unity. • For weak electrolytes, the ionic concentration and ionic strength is small and the error introduced by neglecting activity for concentration is small, i.e., assuming no other salts in solution.

  8. Acid Base Equilibria in Water • CH3COOH + H2O ↔ H3O+ + CH3COOH- • Applying the law of mass action we have • K is the equilibrium constant at a particular concentration also known as the dissociation constant and ionisation constant.

  9. Acid Base Equilibria in Water • If one mole of electrolyte is dissolved in V litres of solution. V = 1/c, where c = concentration in moles/litre. • If the degree of dissociation at equilibrium = α • The amount of unionized electrolyte = 1- α/V moles/litre. • This is also know as Ostwald’s dilution law

  10. Acid Base Equilibria in Water • To be strictly correct As the solution becomes more dilute, the degree of dissociation increases. At infinite dilution the weak acid or base would be totally dissociated.

  11. Strengths of Acids and Bases • Bronsted acids and bases A1-B1 and A2-B2 are conjugate acid base pairs K depends on temperature and the nature of the solvent It is usual to refer to acid base strength of the solvent In water the acid-base pair is H3O+-H2O The conc. of water equals 55.5 moles/litre

  12. Strengths of Acids and Bases If A is an anion acid such as H2PO4- i.e. the second dissociation constant for phosphoric acid • H2PO4- + H2O ↔ HPO42- + H3O+ • NH4+ + H2O ↔ NH3 + H3O+ If A is a cation acid, e.g. ammonium ion. [NH3] = total conc. of ammonia i.e. free NH3 plus NH4OH The H2O is a base since it is accepting a H+

  13. Strengths of Acids and Bases • NH3 + H2O ↔ NH4+ For a Bronsted base, again leaving out H2O In this case the H2O is an acid since it is donating a proton (H+) Since Kw = [H+][OH-] A large pKa corresponds to a weak acid and a strong base

  14. Strengths of Acids and Bases For very weak or slightly ionized electrolytes, the relationship can be reduced since α may be neglected in comparison to unity For any two weak acids or bases at a given dilution V (in litres) we have

  15. Strengths of Acids and Bases

  16. Strengths of Acids and Bases Data expressed as acidic dissociation constants The basic dissociation constant may be obtained from the relationship pKa (acidic) + pKb (base) = Kw (water) = 10-14 @ 25oC

  17. Strengths of Acids and Bases • Consider the reactions: • H2S ↔ HS- + H+ • HS- ↔ S2- + H+

  18. Strengths of Acids and Bases • E.g.: A saturated aqueous solution of H2S is approximately 0.1 M. Both the equilibrium equations must be satisfied simultaneously

  19. Strengths of Acids and Bases • By substituting the values for [H+] and [HS-] into: Which is the value of K2

  20. Le Chatelier's Principle • In 1884 the French chemist and engineer Henry-Louis Le Chatelier proposed one of the central concepts of chemical equilibria. Le Chatelier's principle can be stated as follows: • A change in one of the variables that describe a system at equilibrium produces a shift in the position of the equilibrium that counteracts the effect of this change. • If a chemical system at equilibrium experiences a change in concentration, temperature, volume, or total pressure, then the equilibrium shifts to counter-act the imposed change.

  21. Common Ion Effect • Remember: H2S ↔ HS- + H+ • HS- ↔ S2- + H+

  22. Common Ion Effect • The concentration of an ion in solution may be increased by the addition of another compound that produces the same ion on dissociation. • E.g. The S2- ion conc. by addition of 0.25 M HCl Thus by addition of 0,25 M H+ the sulphide concentration is reduced from 1 x 10-15 to 1.7 x 10-22

  23. Common Ion Effect • Consider the equilibrium reaction of acetic acid • CH3COOH ↔ CH3COO- + H+

  24. Common Ion Effect • Effect of addition of 0.1 moles NaAc (8.2 g) to 1000 mL of 0.1 M HAc. Consider the acetic acid first. • 1 – α ≈ 1 • Hence [H+] = 0.00135, [CH3COO-] = 0.00135, and [CH3COOH] = 0.0986

  25. Common Ion Effect • The concentration of sodium and acetate ions produced by addition of the completely dissociated sodium acetate are: • [Na+] = 0.1, and [CH3COO-] = 0.1 mole/litre • The CH3COO- will tend to decrease the ionisation of the acetic acid, since K is constant, and the acetate ion conc. derived from it. • Hence we may write [CH3COO-] = 0.1 • α’ is the new degree of ionisation • [H+] = α’c = 0.1 α’, and [CH3COOH] = (1 – α’)c = 0.1 since α’ is negligibly small.

  26. Common Ion Effect • Substituting in the mass-action equation • The addition of 0.1 M NaAc to 0.1M acetic acid has decreased the degree of ionisation from 1.35 to 0.018%, and the [H+] from 0.00135 to 0.000018

  27. Solubility Product • For sparingly soluble salts <0.01 M • AgCl (solid) ↔ Ag+ + Cl- • The velocity of the reactions depends on temperature

  28. Solubility Product • v1 = k1 • v2 =k2[Ag+][Cl-] • At equilibrium k1 =k2[Ag+][Cl-] • [Ag+][Cl-] = k1/k2 =SAgCl • Again to be strictly correct activities and not concentrations should be used. At low concentration the activities are practically equal to concentration.

  29. Solubility Product

  30. Solubility Product – Inert Electrolyte • In the presence of moderate concentrations of salts, the ionic strength will increase. This will, in general lower the activity coefficient of both ions, and consequently the ionic concentrations and (and therefore the solubility) must increase in order to maintain the solubility product constant. • E.g. fA+decreased from 1 – 0.8, the activity will decrease and the concentration will increase in order to maintain the correct activity conc.

  31. Solubility Product • The solubility increases by the addition of electrolytes with no common ions

  32. Solubility Product – Effect of Acids • M+ + A- + H+ + Cl- ↔ HA + M+ + Cl- • If the dissociation constant of the acid HA is small, the anion A- will be removed from the solution to form the un-dissociated acid HA. Consequently more of the solid will pass into solution to replace the anions removed and this process will continue until equilibrium is established [M+][A-] = SMA • Fe2+ + CO32- ↔ Fe2CO3↓ • kSFe2CO3 = [Fe2+][CO32-] • H2CO3 ↔ H+ + HCO3- K1 = 4.3 x 10-7

  33. Solubility Product – Effect of Acids • HCO3- ↔ H+ + CO3- K2 = 5.6 x 10-11 • CO32- + H+ → HCO3- • Also for sparingly soluble sulphates, Ba, Sr and Pb Ba2+ + SO42- + H+ + Cl- ↔ HSO4- + Ba2+ + Cl- Since the K2 is comparatively large HSO4- ↔ H+ + SO42- (pKa = 1.92), the effect of addition of a strong acid is relatively small.

  34. Complex Ions • The increase in solubility of a precipitate upon the addition of excess of the precipitating agent is frequently due to the formation of a complex ion. • E.g. the ppt of silver cyanide • SAgCN = [Ag+][CN-] because the solubility product is exceeded • K+ + CN- + Ag+ + NO3- → AgCN↓ + K+ + NO3- or Ag+ + CN- → AgCN↓ The ppt dissolves on addition of excess potassium cyanide due to the formation of the complex ion [Ag(CN)2]- AgCN (solid) + CN- (excess) ↔ [Ag(CN)2]- a soluble complex ion. K[Ag(CN)2] a soluble complex salt.

  35. Instability Constants of Complex Ions The complex ion formation renders the concentration of the silver ion concentration so small that the solubility product of silver cyanide is not exceeded. Also bear in mind that the CN- ion is also in excess.

  36. Instability Constants of Complex Ions • Cu2+ + NH4OH ↔ Cu(OH)2↓ + NH4+ • Cu(OH)2 + 4NH4+ → Cu(NH4)42+ + OH-

More Related