AP®︎ Chemistry: Unit 8 Flashcards
Master key terms and definitions for Unit 8 of AP Chemistry – Acids and Bases to help you prep for quizzes and the AP exam.
Temperature Dependence of Kw and Neutral pH
As temperature changes, Kw changes, so neutral water does not always have pH 7.
As temperature changes, Kw changes, so neutral water does not always have pH 7.
Use [H3O+][OH-] = Kw to solve for the unknown ion concentration.
Water self-ionizes to form H3O+ and OH−, giving Kw = [H3O+][OH−].
In pure water at 25°C, [H3O+] = [OH−] = 1.0 × 10^-7 M.
In aqueous solution, H3O+(aq), often written H+(aq), represents the protonated form of water.
At 25°C, water’s ion product gives pKw = 14.0, so pH + pOH = 14.0.
These acids fully ionize, so hydronium concentration equals the initial acid concentration.
These bases fully dissociate, giving hydroxide equal to one or two times formula concentration.
pH is the negative base-10 logarithm of hydronium ion concentration: pH = -log[H3O+].
pOH is the negative base-10 logarithm of hydroxide ion concentration: pOH = -log[OH−].
The common strong acids are HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4 (strong for its first proton); they dissociate essentially completely in water.
A single-proton acid partially transfers H+ to water, forming H3O+ and its conjugate base.
A weak base partially accepts H+ from water, forming OH- and its conjugate acid.
At equilibrium, [H3O+] is much smaller than initial acid concentration, and most HA remains un-ionized.
At equilibrium, [OH-] is much smaller than initial base concentration, and most base remains un-ionized.
Use an ICE table with Ka to find [H3O+], then calculate pH = -log[H3O+].
Track initial, change, and equilibrium concentrations, then substitute equilibrium values into Ka or Kb.
If ionization is very small, subtracting x from the initial concentration is treated as negligible.
It relates weak acid equilibrium concentrations, and its negative logarithm indicates acid strength.
It relates weak base equilibrium concentrations, and its negative logarithm indicates base strength.
It is the percentage of initial acid or base molecules that ionize at equilibrium.
It connects hydronium, hydroxide, and conjugate pair constants through water’s ionization equilibrium.
Diluting a weak acid or base increases its percent ionization at equilibrium.
Acids increase [H3O+] in water and bases increase [OH−] in water; strong acids and bases ionize essentially completely, while weak acids and bases ionize only partially at equilibrium.
A conjugate acid-base pair consists of two species that differ by one proton, such as HA/A− or B/HB+; buffers contain such a pair.
Find leftover H+ or OH- by stoichiometry, divide by total volume, then calculate pH or pOH.
Mixtures are analyzed by whether acid remains, base remains, or neither remains after stoichiometric reaction.
When weak acid remains after reaction, the solution contains HA and A- and acts as a buffer.
When weak base remains after reaction, the solution contains B and HB+ and acts as a buffer.
pH = pKa + log([B]/[HB+]), equivalent to pH = pKa - log([HB+]/[B]).
K = Ka × Kb / Kw for HA + B ⇌ A- + HB+.
Compare Ka and Kb: larger Ka gives acidic, larger Kb gives basic, similar values give near-neutral.
A quantitative reaction where H+ and OH- form water, with other ions as spectators.
A product-favored equilibrium where HA reacts with OH− to form A− and water.
A product-favored equilibrium where B reacts with H3O+ to form HB+ and water.
An equilibrium reaction where HA and B form A− and HB+ without going to completion.
At equivalence, A− hydrolyzes with water to produce OH−, making the solution slightly basic.
At equivalence, HB+ hydrolyzes with water to produce H3O+, making the solution slightly acidic.
In a buffer, the conjugate base consumes added H+ and the weak acid consumes added OH−, resisting large pH changes.
A buffer is a solution containing a weak acid and its conjugate base or a weak base and its conjugate acid that resists large pH changes.
For a buffer made from a weak acid and its conjugate base, pH = pKa + log([A−]/[HA]).
A controlled neutralization used to determine an unknown acid or base concentration.
Analyte is the unknown solution; titrant is the known solution added from a buret.
A graph of pH versus volume of titrant added during a titration.
The point where stoichiometric amounts of acid and base have completely reacted.
Use balanced-reaction mole ratios at equivalence to relate titrant moles to original analyte moles.
The dominant particles after neutralization determine pH through reactions with water.
Starts at extreme pH, changes gradually, then has a steep jump centered at pH 7.
Both show a buffer region and half-equivalence point, but equivalence pH is basic or acidic.
At equivalence, pH depends on whether the remaining species reacts with water.
Multiple equivalence points reveal proton count, major species, and pKa values for each step.
At the half-equivalence point of a weak acid or weak base titration, conjugate pair concentrations are equal, so pH = pKa for an acid titration or pOH = pKb for a base titration.
Stronger acids have weaker, more stable conjugate bases; stronger bases have weaker conjugate acids.
Hydrogens bonded to electronegative atoms or in acidic groups are most likely to transfer as H+.
Weaker H-X bonds give stronger acids; for hydrogen halides, acidity increases down the group.
More electronegative atoms stabilize negative charge in the conjugate base, increasing acid strength.
Electron-withdrawing atoms pull electron density through sigma bonds, stabilizing nearby negative charge.
Delocalizing negative charge over multiple atoms stabilizes the conjugate base and strengthens the acid.
Stronger oxyacids have more electron-withdrawing central atoms or higher oxidation states, weakening the O-H bond.
Organic acids containing a COOH group that are weak acids because their conjugate bases are resonance-stabilized.
COOH compounds are more acidic than ROH compounds because carboxylate ions are resonance-stabilized but alkoxides are not.
Alkali metal hydroxides and Ca(OH)2, Sr(OH)2, Ba(OH)2 dissociate completely to produce OH-.
Resonance-stabilized conjugate bases of carboxylic acids that are weak bases and can accept H+.
Alkyl groups increase amine basicity by donating electron density to nitrogen's lone pair.
Ammonia, amines, and carboxylate ions are weak bases that only partially react with water.
A quantity written as the negative base-10 logarithm of a value, such as pH or pKa.
For a conjugate acid-base pair in water at 25°C, their values add to 14.
If pH < pKa, HA predominates; if pH > pKa, A− predominates; if equal, concentrations match.
Compare solution pH to the pKa of its conjugate acid to predict protonated versus unprotonated form.
A buffer works best when solution pH is within about 1 unit of the pKa.
Weak acids or bases whose protonated and deprotonated forms have different colors or properties.
Color change occurs mainly over about pKa ± 1, where both indicator forms are present.
Select one whose pKa or transition range is close to the equivalence-point pH.
When conjugate acid and base concentrations are equal, pH = pKa; neither form predominates and buffering is strongest.
pKa = -log(Ka); lower pKa values correspond to larger Ka values and stronger acids.
Their conjugates are negligible bases or acids, so they cannot neutralize added H+ or OH-.
Both conjugate pair members must be present in significant, similar amounts for effective pH resistance.
A buffer contains a conjugate acid-base pair: a weak acid with its conjugate base or a weak base with its conjugate acid.
A buffer is most effective when conjugate acid and base concentrations are equal.
Small added acid or base changes [A−]/[HA] only slightly, so pH changes very little.
Use stoichiometry first to find remaining weak acid and formed conjugate base, then apply the equation.
If both species are in the same total volume, their mole ratio equals their concentration ratio.
Higher concentrations of both buffer components increase resistance to pH change without changing pH if their ratio stays constant.
If conjugate acid and base concentrations change proportionally, the pH stays the same.
A buffer neutralizes added acid or base better depending on which conjugate component is present in greater amount.
Buffer capacity is the amount of added acid or base a buffer can neutralize before its pH changes significantly.
A buffer neutralizes added acid or base more effectively in the direction of the component present in greater amount; overall buffer capacity is the amount of acid or base it can absorb before pH changes significantly.
Solubility changes with pH when a salt contains a weak acid/base ion or OH−.
Lower pH increases solubility when the anion is the conjugate base of a weak acid.
Higher pH increases solubility when the cation is the conjugate acid of a weak base.
Removing a dissolved acidic or basic ion shifts dissolution right; adding it shifts dissolution left.
Conjugates of weak acids or bases react with H+ or OH−; conjugates of strong acids or bases are insignificant.
Salts made from insignificant ions have solubilities largely unaffected by pH alone.
Lower pH increases hydroxide salt solubility, while higher pH decreases it by common-ion effect.
