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AP®︎/College Chemistry
Buffer capacity
AP.Chem:
SAP‑10 (EU)
, SAP‑10.D (LO)
, SAP‑10.D.1 (EK)
, SAP‑10.D.2 (EK)
The capacity of a buffer to neutralize added acid or base depends on the concentrations of HA and A⁻ in solution. For a given ratio of [HA] to [A⁻], the greater the concentrations, the higher the overall buffer capacity. When [HA] is greater than [A⁻], the capacity is higher for added base than acid. When [A⁻] is greater than [HA], the capacity is higher for added acid than base. Created by Jay.
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how would you go about determining the capacity of a basic buffer?(2 votes)
Buffer capacity is defined as the number of moles of acid or base that have to be added to 1 liter to cause its pH to change by 1 unit.
As a formula this is: B = n/ΔpH, where B is buffer capacity (which is unitless), n is the number of moles of acid or base that was added to the buffer per liter of the buffer, and ΔpH is the pH difference between the initial buffer’s pH and the pH after the addition of acid or base to the buffer.
A larger B value would mean the buffer has the capacity to neutralize larger amounts of acid or base before it no longer functions.
Hope that helps.(2 votes)
where sis 4.70 com from when doing ph(1 vote)- That's the pKa of acetic acid, or the -log(Ka), which is a measure of how strong an acid is. A pKa is a constant specific to a chemical and can be found in most chemistry textbook's data sections. In a problem/exam situation they're almost guaranteed to provide this information.
Hope that helps.(1 vote)
Video transcript
- [Instructor] Buffer capacity
refers to the amount of acid or base a buffer can neutralize
before the pH changes by a large amount. An increased buffer capacity means an increased amount
of acid or base neutralized before the pH changes dramatically. Let's compare two buffers; buffer solution one and
buffer solution two, and see which one has the
higher buffer capacity. Buffer solution one has a concentration of acidic acid of 0.250 molar and a concentration of acetate
anion also 0.250 molar. Buffer solution two also
consists of acidic acid and the acetate anion. However, in this case, both concentrations are 0.0250 molar. So buffer solution one
has a higher concentration of both acidic acid and the acetate anion. Let's calculate the initial
pH of both buffer solutions using the Henderson-Hasselbalch equation. In the Henderson-Hasselbalch equation, the pH of the solution is equal
to the pka of the weak acid, which for both buffers is a acidic acid plus the log of the concentration
of the conjugate base divided by the concentration
of the weak acid. In this case, the conjugate
base is the acetate anion, so the concentration of the acetate anion divided by the concentration
of acidic acid. For buffer solution number one, the concentration of the acetate anion is equal to the
concentration of acidic acid. Therefore, the ratio
of their concentrations is equal to one. And it's the same idea for
buffer solution number two, the concentration of the acetate anion is equal to the
concentration of acidic acid. Therefore, the ratio
of their concentrations is also equal to one for buffer two. Since the ratio of the
concentrations is equal to one, the log of one is equal to zero, and the pka value of acidic
acid at 25 degrees Celsius is equal to 4.74. So the pH of both buffer
solutions is equal to 4.74 plus zero, or just 4.74. So we're starting with
two buffer solutions, each at a pH of 4.74, and
to those buffer solutions, we're gonna add 0.0200
moles of hydroxide anions. And by calculating the pH change after adding the hydroxide anions, we'll be able to see which buffer has the
higher buffer capacity. The added hydroxide
anions will be neutralized by the weak acid that's
present in the buffer system. So that's acidic acid. So acidic acid is going to
react with the hydroxide anions. And to make the math easier, we're gonna assume that the total volume of the buffer solutions
is equal to one that liter both before and after
the addition of the base. So if the total volume of
the solution is one liter and the concentration of acidic acid is equal to 0.250 molar in buffer one, that means there's 0.250 moles
of acidic acid in the buffer. So 0.250 moles of a acidic acid will react with the 0.0200 moles of hydroxide anions that we're adding to buffer solution one. Let's calculate the pH
of buffer solution one after the addition of
the hydroxide anions. The hydroxide anions
react with a acidic acid to form water and the acetate anion. In buffer solution one, the initial moles of both acidic acid and the
acetate anion are both 0.250. And to that buffer solution we're adding 0.0200 moles of hydroxide anions. To help us find the final
moles of acidic acid and the acetate anion, we're
gonna use an ICF table, where I is the initial amount of moles, C is the change in moles and F
as the final amount of moles. For this reaction, hydroxide anions are the limiting reactant. So we're gonna use up all of the 0.0200 moles
of hydroxide anions, and we're left with nothing. Since the mole ratio of
hydroxide anions to acidic acid is a one-to-one mole ratio, we're also gonna use up 0.0200 moles of acidic acid, which leaves us with 0.230
moles of acidic acid. And for the acetate anion,
there's also a coefficient of one in the balanced equation. So for losing 0.0200
moles for the reactants, we're gaining 0.0200 moles
for the acetate anion, which gives us a final amount of moles of the acetate anion of 0.270. Now that we have our final moles, we're ready to calculate the
pH of the buffer solution. So the pH is equal to the pka value, which for a acidic acid is 4.74, plus the log of the ratio
of the concentrations. And we could put in the concentrations, but since concentration or molarity is equal to moles divided by liters, a ratio of the concentrations would just have the volume
cancel out because it's the same for both our conjugate
base and our weak acid. So a ratio of the moles is the same thing as a ratio of the concentrations in the Henderson-Hasselbalch equation. And we can get our moles
directly from our ICF table. So the moles of the acetate
anion are equal to 0.270, and the moles of a acidic
acid is equal to 0.230. And when we solve for the pH, we find the pH of the
solution is equal to 4.81. So buffer one started at a pH of 4.74, and after the addition
of the hydroxide anions, the pH rose a little bit to 4.81, however, that's a relatively
small change in the pH. So buffer one did a pretty good job of resisting a large change in pH. Next, let's calculate the pH of buffer two after the addition of
the hydroxide anions. So the initial moles of a
acidic acid in buffer two is equal to 0.0250 moles, which is the same number of
moles as the acetate anion, so 0.0250, and the hydroxide anions that we add is equal to 0.0200 moles. Once again, hydroxide anions
are the limiting reactant, so all of the hydroxide anions are used up and we're left with zero moles. Since the mole ratio of
hydroxide anions to acidic acid is one to one, we use up the
same amount of acidic acid, 0.0200 moles, and we're left with 0.0050
moles of a acidic acid after the neutralization. And since there's a one as a coefficient in front of the acetate anion, the acetate anion is
going to gain 0.0200 moles for a final moles of 0.0450. Next, we can calculate
the pH of the solution using the Henderson-Hasselbalch equation. So the pH is equal to the pka
value of acidic acid of 4.74, plus the log of the concentration
of the conjugate base divided by the concentration
of the weak acid. And since we can
substitute moles for that, we can grab those from
our ICF table and plug in the moles of acetate anion
and the moles of acidic acid, and then solve for the pH. When we solve for the pH, we find that the pH of buffer
solution two is equal to 5.69. So buffer solution two
started at pH of 4.74, and after the addition
of the hydroxide anions, the pH rose to 5.69. That's a relatively large increase in the pH of the solution. So going back to buffer solution one, the initial pH was 4.74, and the pH rose to 4.81 upon
the addition of the base. For buffer solution
two, we started at 4.74 and the pH rose to 5.69. Therefore buffer solution
one had a higher capacity to neutralize the added base. And so we say that buffer solution one has the higher buffer capacity. And since buffer solution two had the more dramatic change in pH upon the addition of
the same amount of base, buffer two has a decreased
capacity to neutralize the base compared to buffer one. So we say that buffer two has
a decreased buffer capacity. Remember that the only difference
between these two buffers was that buffer one had
a higher concentration of acidic acid and the acetate anion. Therefore the higher the
concentration of the weak acid and the conjugate base, the
higher the buffer capacity. We just looked at two buffer solutions in which the concentrations of weak acid and conjugate base were equal. However, a buffer solution
doesn't have to start with equal concentrations of the weak acid and its conjugate base. For example, when the
concentration of weak acid is greater than the
concentration of conjugate base, the buffer has a higher capacity for added base than added acid. And when the concentration
of conjugate base is greater than the
concentration of weak acid, the buffer has a higher
capacity for added acid than added base. As an example, let's think about blood, which has a pH of 7.4. The major buffer system used
to control the pH of blood is the carbonic acid
bicarbonate buffer system. In blood, the concentration of bicarbonate is greater than the
concentration of carbonic acid. And since bicarbonate
is the conjugate base to carbonic acid, the
concentration of the conjugate base is greater than the
concentration of the weak acid. And therefore the buffer
system has a higher capacity for added acid than added base. The reason why the buffer system needs to have a higher
capacity for added acid is because the products of metabolism that enter the bloodstream
are mostly acidic and therefore the bicarbonate anion can react with those acidic
products and neutralize them. Therefore, the buffer system
is able to resist large changes to the pH of blood.