Applications of aqueous equilibria

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physics & chemistry

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published 28/07/2008

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When AgNO3 is added to a saturated solution of AgCl, it is often described as a source of a common ion, the Ag+ ion. By definition, a common ion is an ion that enters the solution from two different sources. Solutions to which both NaCl and AgCl have been added also contain a common ion; in this case, the Cl- ion. There is an effect of common ions on solubility product equilibria. The common-ion effect can be understood by considering the following question: What happens to the solubility of AgCl when we dissolve this salt in a solution that is already 0.10 M NaCl? As a rule, we can assume that salts dissociate into their ions when they dissolve. A 0.10 M NaCl solution therefore contains 0.10 moles of the Cl- ion per liter of solution. Because the Cl- ion is one of the products of the solubility equilibrium, LeChatelier's principle leads us to expect that AgCl will be even less soluble in an 0.10 M Cl- solution than it is in pure water.

Solution containing weak acid HA and its salt NaA.
Common ion effect.
Common Ion-ion produced by both acid and its salt.
Equilibrium calculations.
Buffered solutions.
1. Preparing a buffer.
2. Calculate the pH of the following buffer solutions.
3. Shortcuts in buffer calculations.
4. Calculations involving buffered solutions containing weak acids.
Summary.
Buffer capacity.
Titration.
1. Weak acids-strong bases titration.
2. Weak bases-strong acids titration.
3. Polyprotic acid-strong base titration.
Acid-Base indicators.
Determination of equivalence point.
How indicators work.
Indicators.
Solubility equilibria and the solubility product.
1. Why do some solids dissolve in water?
2. The solubility product expression.
3. The relationship between Ksp and the solubility of a salt.
The role of the ion product (Qsp) in solubility calculations.