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Aqueous Equilibria

Aqueous Equilibria. By: Chris Via. Common-Ion Effect. C.I.E.- the dissociation of a weak electrolyte by adding to the solution a strong electrolyte that has an ion in common with the weak electrolyte.

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Aqueous Equilibria

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  1. Aqueous Equilibria By: Chris Via

  2. Common-Ion Effect • C.I.E.- the dissociation of a weak electrolyte by adding to the solution a strong electrolyte that has an ion in common with the weak electrolyte. • The addition of this strong electrolyte causes the equilibrium to shift as well as the pH to change.

  3. Example • The dissociation of HC2H3O2 is: • HC2H3O2 H+1 + C2H3O2-1 • but if NaC2H3O2 is added to the solution, the addition of C2H3O2- 1 from NaC2H3O2 causes the equilibrium to shift to the left and also reducing the concentration of H+1, therefore raising the pH of the solution.

  4. Practice Problem • Calculate the fluoride -ion concentration and pH of a solution containing 0.1 mol of HCl and 0.2 mol of HF in 1.0L solution. (Ka of HF = 6.8 x 10-4) • Now you want to set up an ice box of the dissociation of HF.

  5. Continued HF H+ F- Ka=6.8*10-4 = [H+][F-]/[HF]=(.1+x)x/(.2-x)

  6. Solution • In this case we can assume x is too small to make a difference so (0.1+ x) becomes just 0.1 • So we are left with 0.1 = 6.8 x 10- 4 0.2 • x =1.4 x 10-3 M= [F-1] • [H+1]= (0.1+ x) = 0.1M • -log ([H+]) = -log (0.1) = 1 • The pH=1.00 and the [F-1]= 1.4 x 10-3 M

  7. Buffered Solutions • Buffers are solutions that resist changes in pH upon the addition of small amounts of acids or bases. • A key example of a buffer is human blood which will stay around a pH of 7.4 • Buffers resist changes in pH because they contain both an acidic species to neutralize OH-1 ions and a basic species to neutralize H+1 ions.

  8. Buffers • However, it is key that the acid and base species of the buffer do not neutralize each other. • To do this, buffers are often prepared by mixing a weak acid or base with a salt of that acid or base. • Ex, the HC2H3O2 and C2H3O2-1 buffer can be prepared by adding NaC2H3O2 to a solution of HC2H3O2.

  9. Buffer Capacity • Buffer Capacity is the amount of an acid or base the buffer can neutralize before the pH begins to change significantly. • The greater the amounts of conjugate acid-base pair, the more resistant the solution is to change in pH. • Henderson-Hasselbalch equation is used to calculate the pH of buffers: • pH = pKa + log[base]/[acid]

  10. Sample Exercise • What is the pH of a buffer that is 0.12 M lactic acid and .1 M sodium lactate? (Ka for lactic acid = 1.4 x 10-4) • You want to start out with an ice box of the dissociation of lactic acid • lactic acid hydrogen lactate

  11. Problem Continued • Ka = 1.4 x 10-4 = x (0.1+x)/(0.12-x) • Once again in this case x will be too small so it can be ignored. • X=1.7 x 10 -4 = [H+1] • pH= -log (1.7 x 10-4) = 3.77 • With the Henderson-Hasselbalch equation: • pH= pKa + log([base]/[acid]) • pH= 3.85 + (-0.08) =3.77

  12. Titrations • An acid-base titration is when an acid is added to a base or vice versa. • Acid-Base indicators are usually used to determine the equivalence point which is the point at which stoichiometrically equivalent quantities of acid and base have been brought together.

  13. Acid-Base Titrations • When strong acids and bases are mixed the pH changes very dramatically near the equivalence point • a single drop a this point could change the pH by a number of units.

  14. Titrations

  15. Polyprotic Acids • When titrating with Polyprotic acids or bases the substance has multiple equivalence points. • So for example, in a titration of Na2CO3 with HCl. There are two distinct equivalence points on the titration curve.

  16. Solubility-Product Constant • Ksp= Solubility Product Constant • It expresses the degree to which the solid is soluble in water. • The equation for Ksp is Ksp = [ion]a[ion]b • An example is the expression of the Ksp of BaSO4: • Ksp = [Ba2+][SO42-]

  17. Factors that Affect Solubility • They are: • the presence of common ions • the pH of the solution • presence of complexing agents • The presence of common ions in a solution will reduce the solubility and make the equilibrium shift left.

  18. pH and Solubility • The general rule of solubility and pH is: • The solubility of slightly-soluble salts containing basic anions increases as the pH of the solution is lowered. • This is because the OH- ionis insoluble in water while the H+ ion is very soluble, therefore when a basic solution has a low concentration of OH- ions the salt will be easier to dissolve.

  19. Continued • The more basic the anion, the more the solubility is influenced by the pH of the solution. • However, salts with anions of strong acids are unaffected by changes in pH.

  20. Complex Ions • A characteristic of most metal ions is their ability to act as a Lewis acid when interacting with water (Lewis base). • However, when these metal ions interact with Lewis bases other than water, the solubility of that metal salt changes dramatically.

  21. Complex Ion • Complex Ion-when a metal ion (Lewis acid) is bonded together with a Lewis base other than water. • Ex. Ag(NH3)2+, Fe(CN)6-4 • The stability of a complex ion can be judged by the size of the equilibrium constant for its formation. • This is called the formation constant, Kf • the larger Kf the more stable the ion is.

  22. Amphoterism • Amphoteric-a metal hydroxide that is capable of being dissolved in strong acids or strong bases, but not in water because it can act like an acid or a base. • Ex. Al+3, Cr+3, Zn+2, and Sn+2 • However, these metal ions are more accurately expressed as Al(H2O)6+3. • This is a weak acid and as it is added to a strong base it loses protons and eventually forms the neutral and water-soluble Al(H2O)3(OH)3

  23. Precipitation of Ions • The reaction quotient, Q, can with the solubility product constant to determine if a precipitation will occur. • If Q > Ksp-precipitation occurs until Q = Ksp • If Q = Ksp-equilibrium exists (saturated solution) • If Q < Ksp-solid dissolves until Q = Ksp • Q is sometimes referred to as the ion product because there is no denominator.

  24. Selective Precipitation • Selective Precipitation-the separation of ions in an aqueous solution by using a reagent that forms a precipitate with one or more of the ions. • The sulfide ion is widely used to separate metal ions because the solubilities of the sulfide salts span a wide range and are dependent on the pH of the solution.

  25. Continued • An example is a solution containing both Cu+2 and Zn+2 mixed with H2S gas. • CuS ends up forming a precipitate, but ZnS does not. • This is because CuS has a Ksp value of 6 x 10-37 making it less soluble than ZnS which has a Ksp value of 2 x 10-25

  26. Qualitative Analysis • Qualitative Analysis determines the presence or absence of a particular metal ion, whereas quantitative analysis determines how much of a certain substance is present or produced. • There are 5 main groups of metal ions: • Insoluble Chlorides, Acid-insoluble sulfides, base-insoluble sulfides and hydroxides, insoluble phosphates, and the alkali metal ions and NH4+1 • These can be found on page 673 in your textbook

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