the concentrations of hydronium ions and hydroxide ions are often reported as pH and pOH, respectively
the aqueous ion of hydrogen has two ions that are often used interchangeably
hydrogen ion:
:
hydronium ion:
water autoionizes with an equilibrium constant K_w
neutral solution: pH = pOH
e.g. pure water
at 25°C: pK_w = 14.0 so pH = pOH = 7.0
the value of K_w is temperature dependent
pH of pure, neutral water will deviate from 7.0 at temperatures other than 25°C
strong acid: completely ionizes in aqueous solution to produce hydronium ions and the conjugate base of the acid
e.g. HCl, HBr, HI, HClO₄, H₂SO₄, HNO₃
concentration of H₃O⁺ in a strong acid solution is equal to the initial concentration of the strong acid
the pH of the strong acid solution can be easily calculated
strong base: completely dissociates to produce hydroxide ions
e.g. group I and II hydroxides
concentration of OH⁻ in a strong base solution is equal to:
group I hydroxide: initial concentration
group II hydroxide: double the initial concentration
the pOH (and pH) of the strong base solution can be easily calculated
weak acid and base equilibria
weak acid: reacts with water to produce hydronium ions, but only a small percentage of molecules will ionize
concentration of hydronium is much less than the initial concentration of the acid
vast majority of acid molecules remain un-ionized
solution involves equilibrium between un-ionized acid and conjugate base
equilibrium constant: K_a
weak base: reacts with water to produce hydroxide ions, but only a small percentage of molecules will ionize
concentration of hydroxide in the solution does not equal the initial concentration of the base
vast majority of base molecules remain un-ionized
solution involves equilibrium between un-ionized base and conjugate acid
equilibrium constant: K_b
percent ionization: can be calculated from:
pK_a or pK_b and initial concentration
initial concentration and equilibrium concentration of any species in the equilibrium expression
reactions
strong acid and a strong base
pH of resulting solution may be determined from the concentration of excess reagent
weak acid and strong base
weak acid in excess
buffer solution is formed
pH can be determined from Henderson-Hasselbalch equation
strong base in excess
pH can be determined from the moles of excess hydroxide ion and the total volume of solution
equimolar
pH (slightly basic) can be determined from the equilibrium represented by the equation
weak base and strong acid
weak base in excess
pH can be determined from the Henderson-Hasselbalch equation
strong acid in excess
pH can be determined from the moles of excess hydronium ion and the total volume of solution
equimolar
pH (slightly acidic) can be determined from the equilibrium represented by the equation
weak acid and weak base
titration: controlled conditions for an acid-base reaction
titration curve: plots pH against volume of titrant added
equivalence point (for monoprotic acids/bases): number of moles of titrant added is equal to the number of moles of analyte originally present
can be used to obtain the concentration of the analyte
titrations of strong acids/bases and weak acids/bases
pH determined by the major species in solution
strong acid and strong base: neutral
weak acid/base: basic/acidic
conjugate is present at the equivalence point and can undergo proton-transfer reactions with the surrounding water
half-equivalence point: halfway to the equivalence point
for weak acids/bases, equal concentrations of each species in the conjugate acid-base pair
polyprotic acids: titration curves can be used to determine the number of acidic protons
major species present at any point along the curve
pK_a associated with each proton in a weak polyprotic acid
molecular structure
molecular structure can be used to infer:
the protons on a molecule that will participate in acid-base reactions
the relative strengths of the protons
strong acids (e.g. HCl, HBr, HI, HClO₄, H₂SO₄, HNO₃) have very weak conjugate bases, stabilized by:
electronegativity
inductive effects
resonance
carboxylic acids (include R−COOH) are one common class of weak acid
strong bases (e.g. group I and II hydroxides) have very weak conjugate acids
common weak bases
nitrogenous bases (e.g. ammonia)
carboxylate ions
electronegative elements tend to stabilize the conjugate base relative to the conjugate acid, and so increase acid strength
protonation state: relative concentrations of HA and A⁻
can be predicted by comparing pH to pK_a of acid
solution pH < acid pK_a: acid form has higher concentration than base form
solution pH > acid pK_a: base form has higher concentration than acid form
acid-base indicator: substance that exhibits different properties (e.g. color) in protonated versus deprotonated state
responds to pH of solution
should have a pK_a close to the pH at the equivalence point
buffer solution: contains a large concentration of both members in a weak conjugate acid-base pair
a buffer can stabilize pH
conjugate acid reacts with added base
conjugate base reacts with added acid
Henderson-Hasselbalch equation
pH of buffer is related to the pK_a of the acid and the concentration ratio of the conjugate acid-base pair
consequence of the equilibrium expression associated with the dissociation of a weak acid
described by the Henderson-Hasselbalch equation
adding small amounts of acid/base to a buffered solution does not significantly change the ratio of [A-]/[HA] and thus does not significantly change the solution pH
change in pH on addition of acid or base to a buffered solution is much less than it would have been in the absence of the buffer
buffer capacity: the quantity of acid/base that a buffer can neutralize
increasing the concentration of the buffer components (keeping ratio constant) keeps the pH of the buffer the same but increases its capacity
buffer has more conjugate acid than base: greater buffer capacity for addition of added base than acid
buffer has more conjugate base than acid: greater buffer capacity for addition of added acid than base
solubility of a salt is pH sensitive when one of the constituent ions is:
weak acid
weak base
hydroxide ion (OH)
can be understood qualitatively using le Châtelier’s principle
