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# Chapter 16

### Aqueous Ionic Equilibrium

TermDefinition
Buffer A solution containing significant amounts of both a weak acid and its conjugate base(or a weak base and its conjugate acid)that resists pH change by neutralizing acid of added base. Buffers need significant amounts of both weak acid and its conjugate base
Summarizing Buffer Characteristics 1. Buffers resist pH change.. 2. A buffer contains significant amounts of both a weak acid and its conjugate base.. 3. The weak acid neutralizes added base.. 4. The conjugate base neutralizes added acid
Common Ion Effect The tendency for a common ion to decrease the solubility of an ionic compound or to decrease the ionization of a weak acid or a weak base
Henderson-Hasselbalch Equation An equation used to easily calculate the pH of a buffer solution from the initial concentrations of the buffer components, assuming that the "x is small approximation is valid"... pH = pKa + log([base]/[acid])
The "x is small approximation" Applies to Problems in Which Both of the Following are True 1. Th initial concentrations of the acids (and/or bases) are not too dilute.. 2. The equilibrium constant is fairly small
When Calculating the pH of a Buffer After Adding Small Amounts of Acid or Base, Remember 1. Adding small amount of strong acid to a buffer converts a stoichiometric amount of the base to the conjugate acid.. 2. Adding a small amount of strong base to a buffer converts a stoichiometric amount of the acid the the conjugate base..
Buffering Action 1. When an acid is added to a buffer, a stoichiometric amount of the weak base is converted to the conjugate a acid.. 2. When a base is added to a buffer, a stoichiometric amount of the weak acid is converted to the conjugate base
Two Factors that Influence the Effectiveness of a Buffer 1. The relative amounts of the acid and the conjugate base.. 2. The absolute concentration of the acid and the conjugate base
Effective Buffer for The Relative Amounts of the Acid and the Conjugate Base In order for a buffer to be reasonably effective, the relative concentrations of the acid and the conjugate base should not differ by more than a factor of 10 (must have [base]/[acid] ratio in the range of 0.10-10)
Effective Buffer for The Absolute Concentrations of the Acid and Conjugate Base A buffer is most effective (most resistant to pH change) when the concentrations of acid and conjugate base are high. The more dilute the buffer compoenets, the less effective the buffer
1. Buffer Range Lowest pH for effective buffer occurs when the base is one tenth as concentrated as the acid. Highest pH for effective buffer occurs when the base is 10 times as concentrated as the acid. Effective range is one pH unit on either side of pKa
2. Buffer Range We can adjust the relative amounts of acid and conjugate base to achieve an pH within the effective range for a buffering system. The buffer would be most effective at the center of the range because the buffer components would be exactly equal at that pH
Buffer Capacity The amount of acid or base that can be added to a buffer without causing a large change in pH
Factors That Increase Buffer Capacity 1. The buffer capacity increases with increasing absolute concentrations of the buffer components.. 2. Overall buffer capacity increases as the relative concentrations of the buffer componenets become more similar to each other
Acid - Base Titration A laboratory procedure in which a basic (or acidic) solution of unknown concentration is reacted with an acidic (or basic) solution of known concentration, in order to determine the concentration of the unknown
Indicator A dye whose color depends on the pH of the solution it is dissolved in; often used to detect the endpoint of a titration
Equivalence Point The point in a titration at which the added solute completely reacts with the solute present in the solution; for acid - base titrations, the point at which the amount of acid is stoichiometrically equal to the amounts of base in the solution
1. Summarizing the Titration of a Strong Acid with a Strong Base 1. The initial pH is simply the pH of the strong acid solution to be titrated..
2. Summarizing the Titration of a Strong Acid with a Strong Base 2. Before the equivalence point, H30^+ is in excess. Calculate the [H30^+] by subtracting the number of moles of added OH^- from the initial number of moles of H30^+ and dividing by the total volume
3. Summarizing the Titration of a Strong Acid with a Strong Base 3. At the equivalence point, neither reactant is in excess and the pH = 7.00..
4. Summarizing the Titration of a Strong Acid with a Strong Base 4. Beyond the equivalence point, OH^- is in excess. Calculate the [OH^-] by subtracting the initial number of moles of H30^+ from the number of moles of added OH^- and dividing by the total volume
Acid Strength and Titration Equivalence Point The volume at the equivalence point in an acid-base titration does not depend on whether the acid being titrated is strong or weak, but only on the amount in moles of acid present in solution before titration begins and on concentration of added base
1. Summarizing Titration of a Weak Acid with a Strong Base 1. The initial pH is that of the weak acid solution to be titrated. Calculate the pH by working an equilibrium problem using the concentration of the weak acid as the initial concentration
2. Summarizing Titration of a Weak Acid with a Strong Base 2. Between the initial pH and the equivalence point, the solution becomes a buffer. Use the reaction stoichiometry to calculate the amounts of each buffer component and then use the Henderson-Hasselbalch equation to calculate the pH
3. Summarizing Titration of a Weak Acid with a Strong Base 3. Halfway to the equivalence point, the buffer components are exactly equal and pH = pKa
4. Summarizing Titration of a Weak Acid with a Strong Base 4. At the equivalence point, the acid has all been converted into its conjugate base. Calculate the pH by working an equilibrium problem for the ionization of water by the ion acting as a weak base.
5. Summarizing Titration of a Weak Acid with a Strong Base 5. For Step 4, calculate the concentration of the ion acting as a weak base by dividing the number of moles of the ion by the total volume at the equivalence point
6. Summarizing Titration of a Weak Acid with a Strong Base Beyond the equivalence point, OH^- is in excess. Ignore the weak base and calculate the [OH^-] by subtracting the initial number of mole of H3O^+ from the number of moles of added OH^- and dividing by the total volume
Endpoint The point of pH change where an indicator changes color
1. Expression for the Ratio of [In^-]/[HIn], the Concentration of Phenolphthalein over its Conjugate Base [In^-]/[HIn] = 10^(pH - pKa)... 1. pH(pH = pKa)...[In^-][HIn]ratio([In^-][HIn] = 10^0 = 1)...Color of Indicator Solution(Intermediate color)..
2. Expression for the Ratio of [In^-]/[HIn], the Concentration of Phenolphthalein over its Conjugate Base 2. pH(pH = pKa + 1)...[In^-][HIn]ratio([In^-][HIn] = 10^1 = 10)...Color of Indicator Solution(Color of In^-).. 3. pH(pH = pKa - 1)...[In^-][HIn]ratio([In^-][HIn] = 10^-1 = 0.10)...Color of Indicator Solution(Color of HIn)
Solubility Product Constant(Ksp) The equilibrium expression for a chemical equation representing the dissolution of a slightly to moderately soluble ionic compound.. For CaF2(s) <---> Ca^2+(aq) + 2F^-, the solubility product constant is(no solids in expression).. Ksp = [Ca^2+][F^-]^2
Molar Solubility The solubility of a compound in units of moles per liter(mol/L).
The Effect of a Common Ion on Solubility In general, the solubility of an ionic compound is lower in a solution containing a common ion than in pure water
The Effect of pH on Solubility In general, the solubility of an ionic compound with a strongly or weakly basic anion increases with increasing acidity(decreasing pH) because the H3O^+ will neutralize the OH^- ions that dissociate from the ionic compound and drive reaction to the right
1. Summarizing the Relationship of Q and Ksp in Solutions Containing an Ionic Compound 1. If Q < Ksp, solution is unsaturated. More of the solid ionic compound can dissolve in solution.. 2. If Q = Ksp, solution is saturated. The solution is holding equilibrium amount of dissolved ions and additional solid will not dissolve in solution
2. Summarizing the Relationship of Q and Ksp in Solutions Containing an Ionic Compound 3. If Q > Ksp, solution is supersaturated. Under most circumstances, the excess solid will precipitate out of a supersaturated solution
Selective Precipitation A process involving the addition of a reagent to a solution that forms a precipitate with one of the dissolved ions but not the others
Qualitative Analysis A systematic way to determine the ions present in an unknown solution. Involves finding the kind of ions present in the solution. In qualitative analysis, specific ions are precipitated successively by the addition of appropriate reagents
Quantitative Analysis A systematic way to determine the amounts of substances in a solution or mixture. Is concerned with the quantity, or the amounts of substances in a solution or mixture
Complex Ion An ion that contains a central metal ion bound to one or more ligands
Ligand A neutral molecule or an ion that acts as a Lewis base with the central metal ion in a complex ion
Formation Constant(Kf) The equilibrium constant associated with reactions for the formations of complex ions
The Effect of Complex Ion Equilibria Solubility The solubility of an ionic compound containing a metal cation that forms complex ions increases in the presence of Lewis bases that complex with the cation
Complex Ion Formation Normally insoluble AgCl is made soluble by the addition of NH3 which forms a complex ion with Ag^+ and dissolves the AgCl
Solubility of an Amphoteric Hydroxide Because aluminum hydroxide is amphoteric, its solubility is pH dependent. At low pH, formation of Al(H2O)6^+3 drives dissolution. At neutral pH, insoluble Al(OH)3 precipitates out of solution. At high pH. formation of Al(H2O)2(OH)4^- drives dissolution
Created by: TimChemistry2