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Chapter 17 Common Ion effect  

The presence of a common ion suppresses the ionization of a weak acid or a weak base. When a compound containing an ion in common with a dissolved substance is added to a solution at equilibrium, the equilibrium shifts to the left.

Buffer Solutions 

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A solution that contains significant concentrations of both members of a conjugate pair (weak acid/conjugate base or weak base/conjugate acid). o Their amounts must be comparable  Conjugate base must be supplied by a dissolved salt. Resist changes in pH when adding small amounts of acids or bases. Assumptions, o Neglect the weak acid that ionizes because you have the common ion which suppresses the ionization in the buffer solution.

Calculating the pH of a Buffer 

Henderson-Hasselbalch equation

or

o  o

Concentrations must be

Preparing a Buffer Solution with a Specific pH 

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Range of the buffer o where pH = pKa ± 1  This enables us to select the appropriate conjugate pair to prepare a buffer with a specific, desired pH. First, we choose a weak acid whose pKa is close to the desired pH. Next, we substitute the pH and pKa values into Henderson-Hasselbalch equation to obtain the necessary ratio of [conjugate base] / [weak acid]. o This ratio can then be converted to molar quantities for the preparation of the buffer.

Acid-Base Titrations 

Strong Acids and Bases Titration o Completely dissociate o Titrant: solution in the burette o Start with a low pH and end with a high pH  Start with an acid and then turn into a base. o The titration curve of a strong acid–strong base titration has a long, steep region near the equivalence point.  Equivalence point is where equal molar amounts of acid and base have been combined The pH at the equivalence point of a strong acid–strong base titration is 7.00.

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Weak Acid and Strong Base Titrations and WB and SA o Titration curves for weak acid–strong base or weak base–strong acid titrations have a significantly shorter steep region.  The pH at the equivalence point of a weak acid–strong base titration is above 7.00.  The pH at the equivalence point of a weak base–strong acid titration is below 7.00

Solubility Equilibria 

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Ksp solubility product constant o Equilibrium constant that indicates to what extent a slightly soluble ionic compound dissolves in H 2O. o Allows us to make quantitative predictions about how much of given ionic compound will dissolve in water. o The smaller the Ksp the less soluble the compound Units to express solubility o Molar solubility (mol/L) o Solubility (g/L)  Can go from molar solubility to solubility by converting to grams To solve for Ksp o Ice table o Fill in what you know o Figure out what we don’t know Using q to predict when precipitate will form o Q less than Ksp no precipitate will form o Q greater than Ksp, precipitate will form

Factors Affecting Solubility 

Common Ion affect o Calculate the molar solubility of silver chloride in a solution that is 6.5 × 10−3 M in silver nitrate 

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pH o

Which of the following compounds will be more soluble in acidic solution than in water: (a) CuS, (b) AgCl, (c) PbSO4

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(a)

o

(b)

o

(c)

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S2− is the conjugate base of the weak acid HS−. S2− reacts with H3O+ as follows: more soluble in acid



Cl− is the conjugate base of the strong acid HCl. Cl− does not react with H3O+.

 is the conjugate base of the weak acid HSO4 2-. It reacts with H3O+ as follows:  more soluble in acid Complex Ion Formation o Complex ion  An ion containing a central metal cation bonded to one or more molecules or ions. o Formation constant  Kf: The equilibrium constant that indicates to what extent complex-ion formation reactions occur.  The larger the kf the more stable the complex ion

Fractional Precipitation 

The separation of a mixture based upon the components solubilities.

Qualitative analysis 

The determination of the type of ions present in a solution....