ch 10 - Acids and Bases Flashcards
Arrhenius acid
dissociates to form an excess of H+ in solution
Arrhenius base
dissociates to form an excess of OH- in solution
Bronsted-Lowry acid
species that donates a hydrogen ion (H+)
Bronsted-Lowry base
species that accepts hydrogen ion (H+)
conjugate acid-base pairs
Bronsted-Lowry acids and bases occur in pairs becaues the definitions require transfer of a proton from the acid to the base
Lewis acid
an electron pair acceptor
Lewis base
an electron pair donor
other terms for Lewis acid-base chemistry
coordinate covalent bond formation; complex ion formation; nucleophile-electrophile interactions
amphoteric
species that reacts like an acid in a base environment and like a base in an acidic environment
amphprotic
a species that can either gain or lose a proton (Bronsted-Lowry)
anion acid nomenclature
acids formed from anions with names that end in -ide have the prefix hydro- and the ending -ic: F- (fluoride) = HF (Hydrofluoric acid); Cl- (Chloride) = HCl (hydrochloric acid); Br- (bromide) = HBr (hydrobromic acid)
naming oxyacids
oxyacids are acids formed from oxyanions; if anion ends in -ite (less oxygen), acid will end with -ous acid; if it ends in -ate (more oxygen), acid will end with -ic acid and prefixes of names are retained
autoionization
water reacted with itself: H20 (l) + H20 (l) ->
water dissociation constant (K sub w)
Kw = [H3O+][OH-] = 10^-14 at 25 degrees C (298 K); at temps above this, Kw will increase as a result of the endothermic nature of the autoionization reaction
p scale
negative logarithm of the number of items: pH and pOH are prototypical examples
pH
-log[H+] = log (1/[H+])
pOH
-log[OH-] = log (1/[OH-]
pH and pOH for aqueous solutions at 298K
pH + pOH = 14; water at equilibrium and 25 degrees C has a concentration of hydroxide ions (10^-7) = to concentration of hydrogen ions: pH of 7 and pOH of 7
how to multiply logs
log (xy) = log x + log y
shortcut to determine p scale values
if the nonlog value is written in proper scientific notation, it will be in the form n x 10^-m where n = number between 1 and 10: -log(n x 10^-m) = -log (n) - log(10^-m) = m - log(n); n will equal number between 1 and 10 which means log n will be a number between 0 and 1 (closer to 1 = closer to 0; closer to 10 = closer to 1) so p value = about m - 0.n where 0.n represents sliding the decimal point of n one position to the left
strong acids and bases
species that completely dissociate into their component ions in aqueous solutions
common strong acids to know
HCl (hydrochloric acid); HBr (hydrobromic acid); HI (hydroiodic acid); H2SO4 (sulfuric acid); HNO3 (nitric acid); HClO4 (perchloric acid)
common strong bases to know
NaOH (sodium hydroxide); KOH (potassium hydroxide); other soluble hydroxides of Group IA metals
weak monoprotic acid dissociation in water
HA (aq) + H2O (l) ->
acid dissociation constant (Ka) of weak acids
K sub a = ([H3O+][A-])/[HA]; the smaller Ka is the weaker the acid and the less it will dissociate; water is not included; weak acid Ka less than 1.0
base dissociation constant of weak base (Kb)
K sub b = ([B+][OH-])/[BOH] from equation BOH (aq) ->
conjugate acid
acid formed from a base gaining a proton
conjugate base
base formed from an acid losing a proton
products of strong acid + strong base reaction
salt and water; neutral when present in equimolar amounts in reactants
product of reaction between strong acid and weak base
forms a salt but often no water because weak bases often are not hydroxides; cation of the salt is a weak acid and will react with water solvent reforming some of the weak base through hydrolysis; pH below 7
products of weak acid and strong base
pH will be in basic range; salt hydrolyzes, with concurrent formation of hydroxide ions; more hydroxide than hydronium ions
weak bases reacted with weak acids
pH of such a solution depends on relative strengths of the reactants: if Kb is greater than Ka it will be basic; and vice versa
square root of an exponent
that exponent divided by 2.
acid equivalent
equal to one mole of H+ (or more properly H3O+) ions
base equivalent
equal to one mole of OH- ions
polyvalent
acids or bases of which each mole liberates more than one acid or base equivalent (ex. H2SO4 (aq) + H2O (l) -> H3O+ + (HSO4)- and then (HSO4)- (aq) + H2O (l) ->
normality
acidity or basicity of solution depends on concentration of acidic or basic equivalents that can be liberated; so each mole of H3PO4 yields three moles (equivalents) of H3O+. Therefore 2 M H3PO4 solution would be 6 N
gram equivalent weight
the mass of a compound that produces one equivalent (one mole of charge) - if divalent then take the molecular weight of molecule and divide by 2 to get how many grams of that molecule it takes to produce one equivalent of H3O+ if it dissociates completely
titration
a procedure used to determine the concentration of a known reactant in a solution; performed by adding small amounts of solution of known concentration (titrant) to a known volume of a solution of unknown concentration (titrand) until completion of reaction (equivalence point)
equivalence point
in acid-base, reached when number of acid equivalents present in original solution equals the number of base equivalents added or vice-versa; strong acid/strong base will have this point at pH 7 but others won’t necessarily
equation for unknown concentration of titrand
N sub a (V sub a) = N sub b (V sub b); Na and Nb = acid and base normalities; Va and Vb = volumes of acid and base solutions
how to select ideal indicator
find pH of the reaction at equivalence point and then select indicator with pKa value closest to it; must be weaker acid or base than the acid or base being titrated
endpoint
the point at which the indicator changes to its final color; should be a negligible difference from equivalence point that can be corrected for or ignored
multiple equivalence on a graph
indicate that it is a polyvalent titration
half-equivalence point
the center of the buffer region (point between two regions on polyvalent graph); occurs when half of a given species has been protonated or deprotonated
buffer solution
consists of a mixture of weak acid and its salt (which is composed of its conjugate base and a cation) or a mixture of a weak base and its salt (which is composed of its conjugate acid and an anion); examples are acetic acid (CH3COOH) and its salt, sodium acetate (CH3COO-Na+); and ammonia (NH3) and its salt, ammonium chloride (NH4+Cl-)
bicarbonate buffer system
conjugate pair in the plasma of blood: H2CO3/(HCO3)-; carbonic acid and bicarbonate; CO2 (g) + H20 (l) ->
Henderson-Hasselbalch equation for weak acid buffer solution
pH = pKa + log [A-][HA]; [A-] = concentration of conjugate base and [HA] = concentration of weak acid; when conjugate base concentration = weak acid concentration pKa = pH because log (1) = 0; buffering capacity is optimal then
Henderson-Hasselbalch equation for weak base buffer solution
pOH = pKb + log [B+]/[BOH]; [B+] = concentration of conjugate acid and [BOH] = concentration of weak base; when concentrations are equal pKb = pOH and buffering capacity is optimal
buffering capacity
ability to which the system can resist changes in pH - if concentrations of acid and its conjugate base were doubled then capacity would double (meaning resistance to pH change would double) but not actual pH (which would not change); buffering capacity is usually maintained within 1 pH unit of the pKa value
