Acids and Bases Flashcards
dissociate to produce an excess of hydrogen ions (H+) in solution; have H at beginning of formula
Arrhenius acid
dissociate to produce an excess of hydroxide ions (OH-) in solution; have OH at end of formula
Arrhenius base
species that can donate H atoms
Bronsted-Lowry acid
species that accept H atoms
Bronsted-Lowry base
electron pair acceptors
Lewis acid
electron pair donors
Lewis base
species that can behave as an acid (in a basic environment) or a base (in an acidic environment)
amphoteric
amphoteric species that donate or accept a proton (H+ ion), behaving as a Bronsted-Lowry acid or base
amphiprotic
amphiprotic example:
water- can accept a H+ to become H3O+ or lose a H+ to become OH-
process where an amphoteric compound reacts with itself (like water); one water molecule can donate a hydrogen atom to another water molecule to produce the hydronium ion (H3O+) and the hydroxide ion (OH-)
autoionization
water dissociation constant (K(w))
K(w) = [H3O+] [OH-] = 10^-14
true at 298 K, only affected by changes in temperature
logarithmic scale for the concentration of hydrogen (hydronium) ions
pH
pH
pH = -log [H+] = log 1/[H+]
logarithmic scale for the concentration of hydroxide ions
pOH
pOH
pOH = -log [OH-] = log 1/[OH-]
pH + pOH = __
pH + pOH = 14
in aqueous solutions at 298 K
completely dissociate into their component ions in aqueous solution
strong acids and bases
do not completely dissociate in solution and have corresponding dissociation constants (K(a) and K(b))
weak acids and bases
acid dissociation constant (K(a))
K(a) = [H3O+] [A-] / [HA]
base dissociation constant (K(b))
K(b) = [OH-] [B+] / [BOH]
acid formed when a base gains a proton
conjugate acid
base formed when an acid loses a proton
conjugate base
when an acid and base react with each other to form a salt and sometimes water
neutralization reaction
defined as one mole of the species of interest
equivalent
equal to one mole of H+ (H3O+) ions
acid equivalent
equal to one mole of OH- ions
base equivalent
the concentration of acid or base equivalents in solution
normality
acids and bases that can donate or accept multiple electrons/protons (by the Bronsted-Lowry definition)
polyvalent/polyprotic
used to determine the concentration of known reactant in a solution; performed by adding small volumes of a solution of known concentration (titrant) to a known volume of a solution of unknown concentration (titrand) until completion of the reaction is achieved at the equivalence point
titration
titration:
has a known concentration and is added slowly to the titrand to reach the equivalence point
titrant
titration:
has an unknown concentration but a known volume
titrand
titration:
midpoint of the buffering region, in which half of the titrant has been protonated (or deprotonated); thus, [HA] = [A-] and a buffer is formed
half-equivalence point
titration:
is indicated by the steepest slope in a titration curve; it is reached when the number of acid equivalents in the original solution equals the number of base equivalents added, or vice-versa
equivalence point
equivalence point equation:
N(A) V(A) = N(B) V(B)
where:
N(A) and N(B) = acid and base normalities
V(A) and V(B) = volumes of acid and base solutions
have equivalence points at pH = 7
strong acid + strong base titrations
have equivalence points at pH > 7
weak acid + strong base titrations
have equivalence points at pH < 7
weak base + strong acid titrations
can have equivalence points above or below 7, depending on the relative strength of the acid and base
weak acid + weak base titrations
titration:
weak acids or bases that display different colors in their protonated and deprotonated forms; ____ chosen for a titration should have a pK(a) close to the pH of the expected equivalence point
indicators
titration:
when the indicator reaches its final color in a titration
endpoint
titration:
observed in polyvalent/polyprotic acid and base titrations
multiple buffering regions and equivalence points
consists of a weak acid and its conjugate salt or a weak base and its conjugate salt; they resist large fluctuations in pH; e.g. a solution of acetic acid (CH3COOH) and its salt, sodium acetate (CH3COO- Na+), or a solution of ammonia (NH3) and its salt, ammonium chloride (NH4+ Cl-)
buffer solutions
refers to the ability of a buffer to resist changes in pH; maximal ____ is seen within 1 pH point of the pK(a) of the acid in the buffer solution
buffering capacity
quantifies the relationship between pH and pK(a) for weak acids and between pOH and pK(b) for weak bases; when a solution is optimally buffered, pH = pK(a) and pOH = pK(b)
Henderson-Hasselbach equation
Henderson-Hasselbach equation
pH = pK(a) + log [A-] / [HA]
pOH = pK(b) + log [B+] / [BOH]
where:
[A-} = concentration of the conjugate base
[HA] = concentration of the weak acid
[B+] = concentration of the conjugate acid
[BOH] = concentration of the weak base
