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  • list of strong acids (6)

    HNO3 < H2SO4 < HCl < HBr < HI < HCLO4

  • list of strong bases (6)

    LiOH < NaOH < KOH < Sr(OH)2 < Ca(OH)2 < Ba(OH)2

  • (𝛂) more stable base resonance structure

    (𝛂) stronger acid

  • (𝛃) larger atoms (down a group)

    (𝛃) stronger acid

  • (𝛄) electronegative surroundings

    (𝛄) stronger acid

  • group 2 hydroxide

    release 2 OH-, more base per mole

  • larger ions

    separate more easily

  • more soluble

    stronger base

  • arrhenius acid

    substance that produces H+ ions

  • arrhenius base

    substance that produces OH- ions

  • electrolyte (acid/base)

    substance that forms ions in solution

  • electrolyte strength

    depends on extent of dissociation

  • strong acid/base

    complete dissociation

  • weak acid/base

    partial dissociation

  • brønsted-lowry acid

    proton (H+) donor

  • brønsted-lowry base

    proton (H+) acceptor

  • amphoteric substance

    can act as an acid or a base

  • conjugate acid-base pair

    2 substances related by the transfer of a proton

  • conjugate acid

    a base with a proton added to it

  • conjugate base

    an acid with a proton removed from it

  • binary acid

    H-Y

  • oxyacids

    H-O-Y

  • acidic H-Y bond

    H atom has positive dipole moment

  • increasing acidity

    decreasing bond strength

  • affects H-O-Y strength (2)

    • electronegativity of Y

    • number of O atoms bonded to Y

  • far right equilibrium

    complete ionization

  • monoprotic acid

    only 1 ionizable proton

  • hydrofluoric acid

    HF (weak)

  • acetic acid

    HC2H3O2 (weak)

  • formic acid

    HCHO2 (weak)

  • sulfurous acid

    H2SO3 (weak)

  • carbonic acid

    H2CO3 (weak)

  • phosphoric acid

    H3PO4 (weak)

  • acid ionization constant (Ka)

    quantifies acid strength

  • smaller Ka

    less ionization in water

  • pKa

    -log(Ka)

  • autoionization

    water acts as an acid/base with itself

  • ion product constant (Kw)

    [H3O+] x [OH-] = 1.0 x 10^-14

  • acidic solution

    [H3O+] > [OH-]

  • basic solution

    [H3O+] < [OH-]

  • pH

    -log[H3O+]

  • pH scale [H+]

    decreases by a factor of 10

  • smaller pKa

    stronger acid

  • Ka

    [HA] [A-] / [HA]

  • % ionization of a weak acid

    [H3O+] at equilibrium divided by initial [HA]

  • with increasing [HA] initial (2)

    • [H3O+]eq increases

    • % ionization decreases

  • [H3O+] for a mix of acids

    [strong acid] = total [H3O+]

  • most weak bases

    act as a base by accepting a proton from water

  • most strong bases

    group 1A and 2A metal hydroxides

  • group 2A: M(OH)2

    dissociates in a single step

  • amines (most weak bases)

    ammonia with 1+ hydrocarbon groups substituted for 1+ hydrogen atoms, have a nitrogen atom w/ a lone pair acting as a proton acceptor

  • anions

    form basic or neutral solutions

  • cations

    form acidic or neutral solutions

  • anion that's a conj. base of a weak acid

    weak base

  • anion that's a conj. base of a strong acid

    pH neutral

  • Kb of an anion

    use Ka of its conj. acid

  • (cation) counterions of strong bases

    pH neutral

  • (cation) conj. acids of weak bases

    weakly acidic

  • (cation) small highly charged metals

    weakly acidic

  • (salts) cation/anion doesn't act as an acid/base

    neutral solution

  • (salts) anion acts as a base

    basic solution

  • (salts) cation acts as an acid

    acidic solution

  • (salts) cation/anion acts as an acid/base

    ion with the higher K determines pH

  • polyprotic acids

    ionizes in successive steps, Ka1 > Ka2

  • lewis acid

    electron pair acceptor (has an incomplete octet)

  • lewis base

    electron pair donor

  • adduct

    product of a lewis acid/base reaction

  • conj. acid of a weak base (low tendency to accept protons)

    high tendency to donate protons

  • hydrolysis

    ion reacts with water to form H+ or OH-

  • conj. base of a weak acid

    produces OH-

  • conj. acid of a weak base

    produces H3O+

  • small, highly charged metal cations

    produces H3O+

  • conjugates of strong acids

    neutral pH

  • conjugates of strong bases

    neutral pH

  • a buffer is made of

    significant amounts of a weak acid and its conjugate base

  • a buffer resists pH by

    neutralizing the added acid/base

  • strong base added to buffer

    neutralized by weak acid

  • strong acid added to buffer

    neutralized by conj. base

  • common ion effect

    a solution with 2 substances that a share common ion decreases the amount of ionization in a weak acid/base

  • [H3O+] = Ka

    in a solution with equal acid and conj. base concentrations

  • henderson-hasselbalch equation

    pH = pKa + log([base]/[acid])

  • acid added to buffer

    stoich. amount of weak base is converted to conj. acid

  • base added to buffer

    stoich. amount of weak acid is converted to conj. base

  • effectiveness of a buffer depends on (2)

    • relative amounts of acid/base

    • absolute concentrations of acid/base

  • capacity of a buffer

    how much acid/base it can neutralize

  • range of a buffer

    pH range over which a particular acid and its conj. base can be effective

  • a buffer is most resistant to pH change when

    [HA] = [A-]

  • buffer effectiveness decreases as

    difference in HA and A- concentrations increase

  • higher (but equal) [HA] and [A-] causes

    decreased pH change

  • effective range for a buffering system

    pKa = ph ± 1

  • equivalence point

    moles of base is stoichiometrically = to moles of acid

  • determine pH of a salt made of a weak acid + base

    pH = 1/2 (pKw + pKa - pKb)

  • half equivelence point

    pH = pKa

  • pKa > pKb

    basic solution

  • pKa < pKb

    acidic solution

  • pKa = pKb

    neutral solution

  • calculate [H3O+] (strong base and acid titration)

    (initial moles H3O+) - (moles OH- added) / total volume

  • calculate [O-] (strong base and acid titration)

    (moles OH- added) - (initial moles H3O+) / total volume

  • before equivalence point (acid titration)

    excess H3O+

  • after equivalence point (acid titration)

    excess OH-

  • volume at equivalence point depends on (3)

    • initial amount of acid

    • base concentration

    • stoichiometry of the reaction

  • after titrating a weak acid with a strong base

    solution becomes a buffer

  • half equivalence point (after adding strong base)

    equal amounts of weak acid and conjugate base

  • at equivalence point

    solution contains an ion acting as a weak base

  • weak acid + strong base titration

    basic equivalence point

  • polyprotic acid titration

    volume needed to reach equivalence point 1 = additional volume needed to reach equivalence point 2

  • larger Ka

    smaller 1/2 equivalence point

  • larger Kb

    higher pH at 1/2 equivalence point