Ch 18: Aqueous Ionic Equilibrium Flashcards
Define was a buffer is and the most important example of one.
A chemical system that resists pH changes by neutralizing added acid or base.
An important buffer in blood is a mixture of carbonic acid and the bicarbonate ion.
How does low blood pH, which results in acidosis, affect the equilibrium between hemoglobin (Hb) and oxygen?
The excess acid causes the equilibrium to shift to the left, reducing the blood’s ability to carry oxygen.
One treatment for ethylene glycol poisoning is the administration of _____ _______. The two molecules are similar enough that the liver enzyme that catalyzes the metabolism of ethylene glycol also acts on this substance, but the enzyme has a higher affinity for it than for ethylene glycol. Consequently, the enzyme preferentially metabolizes ethyl alcohol, allowing the unmetabolized ethylene glycol to escape through the urine. If administered early, this treatment can save the life of a dog or cat that has consumed ethylene glycol.
ethyl alcohol
A buffer contains either:
significant amounts of a weak acid and its conjugate base
or
significant amounts of a weak base and its conjugate acid
A weak acid by itself, even though it partially ionizes to form some of its conjugate base, does not contain sufficient base to be a buffer. Similarly, a weak base by itself, even though it _________ _______ water to form some of its conjugate acid, does not contain sufficient acid to be a buffer. A buffer must contain ___________ amounts of both a weak acid and its conjugate base (or vice versa).
partially ionizes
significant
As long as the amount of strong base (or strong acid) added to a buffer solution is smaller than the amount of weak acid (or conjugate base present), then the resulting __ ______ will be small.
pH change
Consider a solution that initially contains HC2H3O2 and NaC2H3O2, each at a concentration of 0.100 M. The acetic acid ionizes according to the reaction:
However, the ionization of HC2H3O2 in the solution is suppressed compared to its ionization in a solution that does not initially contain any C2H3O2−, because the presence of C2H3O2− shifts the equilibrium to the left (as we would expect from Le Châtelier’s principle). In other words, the presence of the C2H3O2−(aq) ion causes the acid to ionize even less than it normally would, resulting in a less acidic solution (higher pH). This effect is known as the ______ ___ ______, so named because the solution contains two substances (HC2H3O2 and NaC2H3O2) that share a common ion (C2H3O2- ).
State the Henderson-Hasselbalch equation:
This equation, known as the Henderson–Hasselbalch equation, allows us to quickly calculate the pH of a buffer solution from the initial concentrations of the buffer components as long as….
…the x is small approximation is valid.
After you use the x is small approximation to find the pH of a buffer solution, how do you confirm that this approximation is valid?
Confirm that the x is small approximation is valid by calculating the [H3O+] from the pH. Since [H3O+ is formed by ionization of the acid, the calculated [H3O+] has to be less than 0.05 (or 5%) of the initial concentration of the acid in order for the x is small approximation to be valid.
In general, remember that the x is small approximation applies to problems in which both of the following are true: (a) the initial concentrations of acids (and/or bases) are not too dilute; and (b) the equilibrium constant is fairly small. Although the exact values depend on the details of the problem, for many buffer problems this means that the initial concentrations of acids and conjugate bases should be at least ___-___ times greater than the equilibrium constant (depending on the required accuracy).
102−103
When a strong acid is added to a buffer, a stoichiometric amount of the weak base is converted to the conjugate acid. When a strong base is added to a buffer, a stoichiometric amount of the weak acid is converted to the conjugate base.
Lowest pH for effective buffer occurs when the base is ___-_____ as concentrated as the acid.
Highest pH for effective buffer occurs when the base is ___ _____ as concentrated as the acid.
pH = pKa + log [(base)/(acid)] = pKa +log 10 = pKa + 1
The effective range for a buffering system is….
one-tenth
ten times
…one pH unit on either side of pKa.
Normal blood has a pH of 7.4. Notice that the concentration of the bicarbonate ion is 20 times higher than the concentration of carbonic acid and the pH of the buffer is more than one pH unit away from pKa. Why?
The higher bicarbonate ion concentration in blood makes the buffer capacity of blood greater for acid than for base, which is necessary because the products of metabolism that enter blood are mostly acidic.
What is the equivalence point in an acid-base titration?
The point in the titration when the number of moles of base is stoichiometrically equal to the number of moles of acid.
