Ch. 10: Acids and Bases (Complete) Flashcards
defn: Arrhenius acid vs. base
Arrhenius acid: will dissociate to form an excess of H+ in solution
Arrhenius base: will dissociate to form excess of OH- in solution
what types of substance are Arrhenius acids typically limited to?
aqueous acids and bases
what is the pattern for identifying Arrhenius acids? Give 3 common.
contain H at the beginning of their formula
HCl, HNO3, H2SO4
what is the pattern for identifying Arrhenius bases? Give 3 common.
contain OH at the end of their formula
NaOH, Ca(OH)2, Fe(OH)3
defn: Bronsted-Lowry acid vs. base
Bronsted-Lowry acid = a species that donates hydrogen ions (H+)
Bronsted-Lowry base = a species that accepts hydrogen ions
what is the advantage of the Bronsted-Lowry definition over the Arrhenius definition?
Bronsted-Lowry is not limited to aqueous solutions
will most acid-base chemistry reactions on the MCAT work in accordance with the Arrhenius definition or the Bronsted-Lowry definition?
Bronsted-Lowry
why do Bronsted-Lowry acids and bases always occur in pairs? what are these paired referred to as?
the definitions require the transfer of a proton from the acid to the base
referred to as conjugate acid-base pairs
defn: Lewis acid vs. base
Lewis acid: an electron pair acceptor
Lewis base: an electron pair donor
what type of electron pair is transferred in Lewis acid base chemistry?
a lone pair that is not involved in any other bonds
mnemonic: Bronsted Lowry vs. Lewis definitions
the brOnsted lOwry definition revolves around prOtOns
the lEwis definition around ElEctrons
what is the main idea behind what is going on in Lewis acid-base chemistry?
one species pushes a lone pair to form a bond with another
what are three alternative names/scenarios that are equivalent to Lewis acid-base chemistry?
- coordinate covalent bond formation
- complex ion formation
- nucleophile-electrophile interactions
which definition of acids and bases is the most inclusive?
lewis definition
why may we see lewis acids a lot in orgo?
lewis acids are often used as catalysts
defn: amphoteric species
one that reacts like an acid in a basic environment and like a base in an acidic environment
defn: amphiprotic
a substance that can either gain or lose a proton
what is the most common example of an amphoteric species on the MCAT?
water!
explain how water acts as an amphoteric species (how does it react with a base? how does it react with an acid?
what four categories of substances are usually considered amphoteric?
- the partially dissociated conjugate base of a polyvalent acid
- the hydroxides of certain metals
- species that can act as both oxidizing and reducing agents
- amino acids that have a zwitterion intermediate with both cationic and anionic character
nomenclature: acids formed from anions with names that end in -ide
prefix: hydro-
ending: -ic
name + nomenclature: acids formed with oxyanions
name: oxyacids
anion ends in -ite (less oxygen) –> acid ends in -ous acid
anion ends in -ate (more oxygen) –> acids ends in -ic acid
why is it so important to understand the behavior of acidic and basic compounds in water?
because many acid-base reactions take place in water, especially on the MCAT
defn + eqn + process: autoionization of water
water can react with itself
one water molecule donates a hydrogen ion to another water molecule to produce the hydronium ion (H3O+) and the hydroxide ion (OH-)
Are H+ and H3O+ the same thing?
Kind of. Many courses depict the hydrogen ion simply as H+, rather than H3O+, but it is important to remember that the proton is never isolated in the solution (it is always attached to water or some other species that has the ability to accept it)
Why is the expression for the autoionization of water in equilibrium?
because it is a reversible reaction
value: Kw (water dissociation constant) for pure water at 298 K
each mole of water that autoionizes produces: how many moles of hydrogen/hydronium ions and hyroxide ions? what does this imply about the concentrations of each?
one mole of each
so the concentrations of hydrogen ions and hydroxide ions are always equal in pure water at equilibrium
so the concentration of each of the ions in pure water at equilibrium at 298 K is 10^-7 M
will the concentrations of the two ions always be equal?
no, they will only be equal when the solution is neutral BUT the product of their respective concentrations will always be 10^-14 when the temperature of the solution is 298 K
what is the one factor that affects Kw, like any other equilibrium constant?
Temperature; otherwise Kw will remain unchanged
what happens to Kw at temperatures above 298 K? why?
it will increase (a direct result of the endothermic nature of the autoionization reaction)
defn: pH and pOH scales
logarithmic expressions of the concentrations of acidic and basic solutions
for the concentrations of hydrogen and hydroxide ions
defn: p scale
the negative logarithm of the number of items
eqn: pH and pOH of a solution
pH = -log[H+] = log 1/[H+]
pOH = -log[OH-] = log 1/[OH-]
what are the pH and pOH of pure water at 298 K?
pH =7
pOH = 7
pH + pOH for aq. solns at 298 K = ?
pH + pOH = 14
is pH < 7 acidic or basic? what about pH > 7? (for aq solns at 298 K)
what does this say about pOH? excess of what kind of ions?
pH < 7 = acidic = pOH > 7 = relative excess of hydrogen ions
pH > 7 = basic = pOH < 7 = relative excess of hydroxide ions
as pH increases, pOH decreases by the same amount
Quick Conversions: concentration has a power of ten, for example [H+] = 10^-3. what does pH equal? what does pOH equal?
pH = 3
pOH = 11
Quick Conversions: concentration has a power of ten, for example Kb = 10^-12. what does pKb equal?
pKb = 12
eqn + how do we get to this mathematically: close approximation of p scale value
p value = (approx) m - 0.n
where 0.n represents sliding the decimal point of n one position to the left (dividing n by 10)
explain the rationale behind using 0.n in the equation above
n is a number between 1 and 10, so its log will be a decimal between 0 and 1 (log 1 = 0, log 10 = 1)
the closer n is to 1, the closer log n will be to 0
the closer n is to 10, the closer log n will be to 1
defn + arrow + char: strong acids and bases
completely dissociate into their component ions in aqueous solutions
single-headed arrows (complete dissociation with no reversibility)
when can we assume that the contribution of OH- and H+ ions from the autoionization of water is negligible? when is the contribution important?
negligible: if the concentration of the acid or base is significantly greater than 10^-7 M
important: if the concentration of acid or base is close to 10^-7 M
what are 6 strong acids commonly encountered on the MCAT? (names + formulas)
- HCl (hydrochloric acid)
- HBr (hydrobroic acid)
- HI (hydroiodic acid)
- H2SO4 (sulfuric acid)
- HNO3 (nitric acid)
- HClO4 (perchloric acid)
what are 3 strong acids commonly encountered on the MCAT?
- NaOH (sodium hydroxide)
- KOH (potassium hydroxide)
- soluble hydroxides of group IA metals
what is the assumption underlying the calculation of the pH and pOH of strong acids and bases?
it assumes complete dissociation of the acid or base in solution
recap: what are we referring to when we discuss acids and bases as strong or week? what are we not referring to?
YES: the chemical behavior of an acid or base with respect to its tendency to dissociate (strong bases completely dissociate in aqueous solutions)
NO: concentrations of acid and base solutions (use the terms concentrated or dilute)
defn: weak acids and bases
only partially dissociate in aqueous solutions
eqn + process: how a weak monoprotic acid, HA, dissociates in water
eqn: acid dissociation constant
what does a small Ka mean?
the smaller the Ka –> the weaker the acid –> the less it will dissociate
eqn + process: how a weak monovalent Arrhenius base, BOH, dissociates in water
undergoes dissociation to yield B+ and OH- in solution
eqn: base dissociation constant
what does a small Kb mean?
smaller Kb –> the weaker the base –> the less it will dissociate
in terms of Kb and Ka, how can we generally characterize as a species as a weak base or a weak acid (number)?
Ka < 1.0 = weak acid
Kb < 1.0 = weak base
why do conjugate acid-base pairs exist?
the Bronsted-Lowry definition of an acid-base reaction is one in which a hydrogen ion is transferred from an acid to a base –> the two always occur in conjugate paris
defn: conjugate acid and conjugate base
conjugate acid = the acid formed when a base gains a proton
conjugate base = the base formed when an acid loses a proton
eqn + implic: relationship between Ka and Kb (conj base) or vice versa
Ka and Kb are inversely related
If Ka is large, then Kb is small and vice-versa
A strong acid will produce a ____ conjugate base
A strong base will produce a ____ conjugate acid
A weak acid or base will produce _____ conjugates
A strong acid will produce a VERY WEAK conjugate base
A strong base will produce a VERY WEAK conjugate acid
A weak acid or base will produce WEAK conjugates
why are the conjugates of a strong acid or base sometimes termed inert?
because it is almost completely unreactive
what is the effect of induction on acid strength?
electronegative elements positioned near an acidic proton increase acid strength by pulling electron density out of the bond holding the acidic proton
this weakens proton bonding, facilitates dissociation
THUS: acids that have EN elements nearer to acidic hydrogens are stronger than those that do not
what is the most common use of acid and base dissociation constants?
to determine the concentration of one of the species in solution at equilibrium
func + eqn + result: neutralization reaction
acids and bases may react with each other to form a salt and often water
result: the salt may precipitate out or remain ionize in solution, depending on its solubility and the amount produced
in general, these reactions go to completion
defn: hydrolysis
salt ions react with water to give back the acid or base
the reverse of a neutralization reaction
what are the 4 combinations of strong and weak acids and bases?
- strong acid + strong base: HCl + NaOH –> NaCl + H2O
- strong acid + weak base: HCl + NH3 –> NH4Cl
- weak acid + strong base: HClO + NaOH –> NaClO + H2O
- weak acid + weak base: HClO + NH3 –> NH4ClO
describe: 1. strong acid + strong base: HCl + NaOH –> NaCl + H2O
- product: equimolar amounts of salt and water
- acid and base neutralize each other, resulting soln is neutral
- ions formed in the reaction will not react with water bc they are inert conjugates
describe: 2. strong acid + weak base: HCl + NH3 –> NH4Cl
- product: salt
- water isn’t formed bc weak bases are often not hydroxides
- the cation of the salt is a weak acid and will react with the water solvent, re-forming some of the weak base through hydrolysis
- the increase in concentration of the hydronium ion causes the system to shift away from autoionization, thereby reducing the concentration of hydroxide ion
- consequently, the concentration of the hydronium ion will be greater than that of the hydroxide ion at equilibrium, and as a result, the pH of the solution will fall below seven, making a slightly acidic solution
describe: 3. weak acid + strong base: HClO + NaOH –> NaClO + H2O
- the pH of the solution at equilibrium will be within the basic range because the salt hydrolyzes, with concurrent formation of hydroxide ions
- the increase in hydroxide ion concentration will cause the system to shift away from autoionization, thereby reducing the concentration of the hydronium ion
- the concentration of the hydroxide ion will be greater than that of the hydronium ion at equilibrium, and as a result, the pH of the solution will rise above 7 (slightly basic solution)
describe: 4. weak acid + weak base: HClO + NH3 –> NH4ClO
pH of the resulting solution depends on the relative strengths of the reactants
how is the relative acidity or basicity of an aqueous solution determined?
by the relative concentrations of acid and base equivalents
defn: acid equivalent
equal to one mole of H+ (or, more properly, H3O+) ions
defn: base equivalent
equal to one mole of OH- ions
defn: polyvalent
each mole of the acid or base liberates more than one acid or base equivalent
defn: polyprotic
defined under the Bronsted Lowry definition
each mole of the acid or base liberates more than one proton?
defn: normality
directly indicates the quantity of the acidic or basic capacity
defn: gram equivalent weight
the mass of a compound that produces one equivalent (one mole of charge)
what are 3 common polyvalent acids?
H2SO4
H3PO4
H2CO3
what are 3 common polyvalent bases?
Al(OH)3
Ca(OH)2
Mg(OH)2
defn: titration + 2 main types
a procedure used to determine the concentration of a known reactant in a solution
- acid-base
- oxidation-reduction
what is the general principle of how a titration is performed?
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
when is the equivalence point reached in acid-base titrations?
when the number of acid equivalents present in the original solution equals the number of base equivalents added (or vice-versa)
does the equivalence point always occur at pH 7?
no, it will be true for a strong acid/strong base titration, but not necessarily otherwise
is there always only one equivalence point?
no, when titrating polyprotic acids or bases ,there are multiple equivalence points, as each acidic or basic conjugate species is titrated separately
at the equivalence point, what is true about the relationship between the number of equivalents of acid and base?
they are equal!
how do you calculate the unknown concentration of the titrand?
Na and Nb are the acid and base normalities
Va and Vb are the volumes of acid and base solutions
make sure the volume use the same units
how is the equivalence point of an acid-base titration determined?
- graphically (plot the pH of the unknown solution as a function of added titrant by using pH meter)
- estimated by watching for a color change of an added indicator
defn: indicator
weak organic acids or bases that have different colors in their protonated and deprotonated states
how does an indicator work?
the binding or release of a proton leads to a change in the absorption spectrum of the molecule, which we perceive as a color change
why must the indicator always be a weaker acid or base than the acid or base being titrated?
the indicator would be titrated first, otherwise!
defn: endpoint (titration)
the point at which the indicator changes to its final color
what is the relationship between the endpoint and the equivalence point in titrations? (if done well)
if the indicator is chosen properly and the titration is done well: the volume difference between the endpoint and the equivalence point is negligible
what are the equivalence point ranges for the three main combos of acid base titrations?
- strong acid + weak base : equivalence point pH < 7
- strong acid + strong base : equivalence point pH = 7
- weak acid + strong base : equivalence point pH > 7
what are the most and lease useful combos of acid base titrations and why?
MOST: involve at least one strong species
LEAST: weak acid/weak base titrations –> not very accurate, rarely performed
–> pH curve lacks the sharp change that normally indicates the equivalence point
–> indicators are less useful bc the pH change is far more gradual
where will the equivalence point be for the different types of titrations?
- strong base titrated into strong acid, pH = 7 equiv pt
- strong base titrated into weak acid, pH > 7 equiv pt
- strong acid titrated into weak base, p < 7 equiv pt
- weak acid and weak base, equiv point near 7
ranges: identifying which type of titration is being shown in a graph by identifying the starting position in the graph
pH»_space; 7 = titrant is strong base
pH > 7 (slightly) = titrant is weak base
pH < 7 (slightly) = titrant is weak acid
pH «_space;7 = titrant is strong acid
then determine where the equivalence point is
analogy: titrations
tug-of-war
the stronger the acid or base, the more it pulls the equivalence point into its pH territory
what is the main differentiating factor between the titration of a polyvalent acid/base titration and a monovalent acid/base titration?
polyvalent titration = multiple equivalence points
defn: half-equivalence point
the center of the buffer region (the point between regions I and II)
it occurs when half of a given species has been protonated or deprotonated
what are the flat regions in the polyvalent titration curves?
buffer regions
what are the three equivalence points in the titrations of acidic and basic amino acids?
- titration of the carboxyl group
- titration of the amino group
- acidic or basic side chain
defn (2) + char (2): buffer solution
a mixture of a 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 (composed of its conjugate acid and an anion)
resist changes in pH when small amounts of acid or base are added
have a narrow range of optimal activity (pKa +/- 1)
what are two common examples of buffers?
- solution of acetic acid (CH3COOH) and its salt sodium acetate (CH3COO-Na+) –> acid buffer
- solution of ammonia (NH3) and its salt ammonium chloride (NH4+Cl-) –> base buffer
defn + char: bicarbonate buffer system
one of the most important buffers in the human body
the H2CO3 (carbonic acid)/HCO3- (conjugate base: bicarbonate) conjugate pair in the plasma part of blood –> forms a weak acid buffer for maintaining the blood’s pH within a narrow physiological range
the majority of the CO2 transported from peripheral tissues to the lungs is dissolved in the plasma in a disguised form through the bicarb buffer system
how do CO2 and water react in the bicarbonate buffer system?
what pH does the bicarbonate buffer system maintain? why does this make sense?
7.4
this is slightly higher than the optimal buffering capacity HOWEVER acidemia is a lot more common than alkalemia, so as acidemia becomes more severe, the buffer system becomes more effective and more resistant to further pH lowering
use + eqn: Henderson-Hasselbalch
used to estimate the pH or pOH of a buffer solution
for weak acid buffer solution:
[A-] = concentration of conjugate base
[HA] = concentration of weak acid
for weak base buffer solution:
[B+] = the concentration of the conjugate acid
[BOH] = concentration of the weak base
what happens with titrations and buffers when [conjugate base] = [weak acid] and when [conjugate acid] = [weak base]?
[conjugate base] = [weak acid]
- pH = pKa
- occurs at half-equivalence points in a titration
- buffering capacity optimal
[conjugate acid] = [weak base]
- pOH = pKb
- buffering capacity optimal
defn: buffering capacity
the ability to which the system can resist changes in pH
generally maintained within 1 pH unit of the pKa value
what happens if the concentrations of both the acid and its conjugate base were doubled (ratio stays the same)?
the pH would not change
the buffering capacity would double (therefore addition of a small amount of acid or base will cause less of a deviation in pH)
