Aqueous equilibria
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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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Aqueous equilibria

Aqueous Equilibria

By: Chris Via

Common ion effect

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.



  • The dissociation of HC2H3O2 is:

  • HC2H3O2 H+ + C2H3O2-

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

Practice problem

Practice Problem

  • Calculate the fluoride -ion concentration and pH of a solution containing .1 mol of HCl and .2 mol of HF in 1.0L solution. (Ka of HF =6.8*10-4)

  • Now you want to set up an ice box of the dissociation of HF.






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



  • In this case we can assume x is too small to make a difference so (.1+x) becomes just .1

  • So we are left with .1x/.2= 6.8*10-4

  • x=1.4*10-3 M= [F-]

  • [H+]= (.1+x)=.1M

  • -log([H+])=-log(.1)=1

  • The pH=1.00 and the [F-]= 1.4*10-3 M

Buffered solutions

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- ions and a basic species to neutralize H+ ions.



  • 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-C2H3O2- buffer can be prepared by adding NaC2H3O2 to a solution of HC2H3O2.

Buffer capacity

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]

Sample exercise

Sample Exercise

  • What is the pH of a buffer that is .12 M lactic acid and .1 M sodium lactate? (Ka for lactic acid = 1.4*10-4)

  • You want to start out with an ice box of the dissociation of lactic acid

  • lactic acid hydrogen lactate

Problem continued

Problem Continued

  • Ka = 1.4*10-4=x(.1+x)/(.12-x)

  • Once again in this case x will be too small so it can be ignored.

  • X=1.7*10-4=[H+]

  • pH= -log(1.7*10-4) = 3.77

  • With the Henderson-Hasselbalch equation:

  • pH= pKa + log([base]/[acid])

  • pH= 3.85 + (-.08) =3.77

Acid base titrations

Acid-Base 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.

  • 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.



Polyprotic acids

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.

Solubility product constant

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-]

Factors that affect solubility

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.

Ph and solubility

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.



  • 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.

Complex ions

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.

Complex ion

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.



  • 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

Precipitation of ions

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.

Selective precipitation

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.



  • 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*10-37 making it less soluble than ZnS which has a Ksp value of 2*10-25

Qualitative analysis

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+

  • These can be found on page 673 in your textbook

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