Aqueous Equilibria

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# Aqueous Equilibria - PowerPoint PPT Presentation

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

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.
• The addition of this strong electrolyte causes the equilibrium to shift as well as the pH to change.
Example
• 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
• 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.
Continued

HF

H+

F-

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

Solution
• 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
• 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.
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-C2H3O2- buffer can be prepared by adding NaC2H3O2 to a solution of HC2H3O2.
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
• 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
• 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
• 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
• 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
• 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
• 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
• 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.
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.
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-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.
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
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-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.
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*10-37 making it less soluble than ZnS which has a Ksp value of 2*10-25
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