Chemistry - Ch 18 - Aqueous Ionic Equilibrium

mainly buffers covers 18.1-18.3 so far

buffer

buffer

a solution containing significant amounts of both a weak acid and its conjugate base (or a weak base and its conjugate acid)

resists pH change by neutralizing added acid or added base

(acid neutralizes added base or base neutralizes added acid)

e.g. HC2H3O2 (acetic acid) and NaC2H3O2 (sodium acetate; conjugate base)

common ion effect

tendency for a common ion to decrease the solubility of an ionic compound or to decrease the ionization of a weak acid or weak base

Henderson-Hasselbalch equation

used to calculate pH of a buffer solution

relates the pH of a buffer solution to the initial concentration of the buffer components

(simplifies the calculation of a buffer solution)

can be used as long as the x is small approximation is valid (because equilibrium concentrations of HA and A- will be essentially identical to initial concentrations, thus [H3O+]=Ka)

note: x is small approximation is valid because so little of the weak acid ionizes compared to its initial concentration; [H3O+]=Ka as long as weak acid and conjugate base concentrations are equal in the buffer

approximate pH of a buffer solution (without calculation)

if [acid]=[base]: pH=pKa

if [acid]>[base]: pH<pKa

if [acid]<[base]: pH>pKa

pH = pKa

if concentration of acid equals concentration of conjugate base

pH > pKa

if concentration of acid is less than concentration of conjugate base

pH < pKa

if concentration of acid is greater than concentration of conjugate base

2 parts of calculating change in pH

stoichiometry calculation

equilibrium calculation

stoichiometry calculation

first part of calculating pH change when an acid or base is added to a buffer

calculation of how the addition changes the relative amounts of acid and conjugate base

if strong acid is added to buffer and neutralized, an amount of the weak base is converted into its conjugate acid

if strong base is added to buffer and neutralized, an amount of the weak acid is converted to its conjugate base

equilibrium calculation

calculation of pH based on the new amounts of acid and conjugate base after the stoichiometry calculation

use ICE table based on a balanced equation for the ionization of the acid

most effective buffer

most resistant to pH changes

when concentrations of acid and conjugate base are equal

buffers become less effective as the difference in relative amounts of acid and conjugate base increases

effective buffer must have a [base]/[acid] ratio in the range of 0.1 to 10

effective range for a buffering system is pKa±1

buffer capacity

the amount of acid or base that can be added to a buffer without causing a large change in pH

increases with increasing absolute concentrations of the buffer components

overall buffer capacity increases as the relative concentrations of the buffer components become more similar (ratio of buffer components gets closer to 1)

acid-base titration

procedure in which a basic (or acidic) solution of unknown concentration reacts with an acidic (or basic) solution of known concentration in order to determine the concentration of the unknown

pH is monitored with either a pH meter or indicator

as acid and base combine, they neutralize each other

indicator

a substance whose color depends on the pH

equivalence point

the point in titration when the number of moles of base is stoichiometrically equal to the number of moles of acid

when titration is complete

neither reactant is in excess, the number of moles of the reactants are related by the reaction stoichiometry

titration curve (or pH curve)

a plot of the pH of the solution during a titration

titration of a strong acid with a strong base

solution starts with containing only strong acid (e.g. HCl)

NaOH is added

initial pH is pH of the strong acid solution

before equivalence point, calculate [H3O+] by subtracting the number of moles of added OH- from the initial moles of H3O+ and dividing by the total volume

pH will always be 7.0 at equivalence point (at 25*C)

beyond equivalence point, calculate [OH-] by subtracting the initial number of moles of H3O+ from the moles added of OH- and dividing by the total volume

titration of a strong base with a strong acid

solution starts with strong base

acid is added, pH decreases

pH at equivalence point is 7.0

(similar to titration of strong acid with strong base, but flipped)

titration of a weak acid with a strong base

begin with weak acid, add strong base until equivalence point is reached

find initial moles of acid

determine number of moles of NaOH needed to neutralize acid

find volume of the base using molarity

solve using Ka and an ICE table

solve for equivalence point

titration of a weak base with a strong acid

similar to calculating points along the curve for the titration of a weak acid with a strong base

solution starts basic and has an acidic equivalence point

titration of a polyprotic acid

pH curve has multiple equivalence points (e.g. titration of a diprotic acid with strong base will result in a pH curve with 2 equivalence points)

end point

the point in titration where an indicator changes color

used to determine the equivalence point

solubility product constant

the equilibrium constant for a chemical equation representing the dissolution of an ionic compound

usually = products multiplied together and raised to the power of coefficients (since solids are omitted and usually reactant is a solid)

can be used to calculate molar solubility of a compound

effect of a common ion on solubility

when a compound with a common ion is added to another solution, solubility either decreases or increases

the solubility of an ionic compound is lower in a solution containing a common ion than in pure water2

effect of pH on solubility

in general, the solubility of an ionic with a basic anion increases with increasing acidity (decreasing pH)

unsaturated solution

a solution that contains less than the equilibrium amount of dissolved ions (when all solid has already been dissolved)

reaction proceeds to the right

Q < Ksp

saturated solution

a solution whose reaction is at equilibrium

Q = Ksp

supersaturated solution

a solution that contains more solute than solvent

forms a precipitate when sufficiently disturbed

Q > Ksp

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

requires a difference in Ksp by a factor of at least 10³

qualitative analysis

a systematic way to determine the ions present in an unknown solution

selective precipitation can be used for it

in the past, used to determine the metals present in a sample

quantitative analysis

a systematic way to determine the amounts of substances in a solution or mixture

complex ion

an ion containing a central metal ion bound to one or more ligands

e.g. Ag(H2O)2+

ligand

a neutral molecule or ion that acts as a Lewis base with the central metal ion (reminder: Lewis acids donate an electron pair)

e.g. H2O in Ag(H2O)2+

will be replaced if a stronger Lewis base is put into a solution containing the subject ion (e.g. NH3)

formation constant

the equilibrium constant associated with the reaction for the formation of a complex ion

calculated the same as any other K value

effect of complex ion equilibria on solubility

the solubility of an ionic compound containing a metal ion that forms complex ions increases in the presence of Lewis bases that complex with the cation