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Buffer Effectiveness, Titrations & pH curves
Section 16.3-16.4
Buffer effectiveness
Buffer effectiveness refers to the ability of a buffer to resist pH change
Effective buffers only neutralize small to moderate amounts of acid or base
Factors influencing buffer effectiveness:
● Ratio of the buffers acid-conjugate base concentrations○ Composition of buffer solution
● Overall absolute concentration○ Concentration of buffer solution
Optimal conditions for buffer effectiveness can be derived from the Henderson-Hasselbalch equation
Reference Table
Relative amounts of Acid and Base
Buffers are most effective when their acid and conjugate base concentrations are equal
Buffers become less effective as the difference between the concentrations of acid and conjugate base increase
In order for a buffer to be effective the difference in concentration should not differ by more than a factor of 10
The math
This can be illustrated by examining two solutions of a generic buffer, pKa=5.00, neutralizing 0.01 mol NaOH.
Both solutions have a volume of 1.0 liter and 0.20 mol total acid and conjugate base
Solution I has equal concentrations of acid and conjugate base while Solution II has 0.18 mol acid (HA) and 0.02 mol base (A-).
Solution I Solution II
OH- + HA ⇒ H2O + A-
0.00 0.10 mol 0.10mol
0.01 mol — —
0.00 mol 0.09 mol .110 mol
OH- + HA ⇒ H2O + A-
0.00 0.18 mol 0.02 mol
0.01 mol — —
0.00 mol 0.17 mol .03 mol
pH =pKa + log ( [base] / [acid] ) =5.00 + log ( 0.1100 / .090 ) =5.09% change=(5.09 - 5.00) / 5.00 X100%
=1.8%
pH =pKa + log ( [base] / [acid] ) =5.00 + log ( 0.03 / .17 ) =4.25% change=(4.25 - 4.05) / 4.05 X100%
=5.0%
Initial pH =5.00 + log (1.00) = 5.00 Initial pH =5.00 + log (0.02/0.18) = 4.05
Reference Table
Before
Addition
After
Before
Addition
After
In the context of the H.-H. equation
pH=pKa + log ( [base] / [acid] )
[base] = [acid] ⇒ [base] / [acid] = 1
log(1)=0
pH=pKa + 0pH=pKa
Concentrations of Acid and Base
Buffers are most effective at high concentrations of acid and conjugate base
The more dilute the buffer components the less effective
Solution I Solution II
OH- + HA ⇒ H2O + A-
0.00 0.50 mol 0.50mol
0.01 mol — —
0.00 mol 0.49 mol .51 mol
OH- + HA ⇒ H2O + A-
Before
Addition
After
0.00 0.05 mol 0.05 mol
0.01 mol — —
0.00 mol 0.04 mol .06 mol
pH =pKa + log ( [base] / [acid] ) =5.00 + log ( 0.51 / 0.49 ) =5.02% change=(5.02 - 5.00) / 5.00 X100%
=0.4%
pH =pKa + log ( [base] / [acid] ) =5.00 + log ( 0.06 / 0.04 ) =5.18% change=(5.18 - 5.00) / 5.00 X100%
=3.6%
Initial pH =5.00 + log (0.50/0.50) = 5.00 Initial pH =5.00 + log (0.050/0.050) = 5.00
Before
Addition
After
Buffer Range
Because a buffer should not differ by more than a factor of 10, we can use the Henderson-Hasselbalch equation to find a pH range:
pH = pKa + log ( [base] / [acid] )
= pKa + log 0.10
= pKa - 1
⇒ the effective pH range for a buffer solution is pKa ± 1
When making a buffer solution, use the pKa then adjust acid-conjugate base ratio
pH = pKa + log ( [base] / [acid] )
= pKa + log 10
= pKa + 1
A: Which acid would one use (combined with its sodium salt) to make a solution buffered at a pH of 4.25?
a.) HClO2 pKa = 1.95b.) HNO2 pKa = 3.34c.) HCHO2 pKa = 3.75d.)HClO pKa = 7.54
Practice
Practice
B: Calculate the ratio of the conjugate base to the acid in formic acid (HCHO2) required to attain the desired pH of 4.25. (pKa = 3.74)
pH = pKa + log ( [base] / [acid] )
4.25 = 3.74 + log ( [base] / [acid] )
4.25 - 3.74 = log ( [base] / [acid] )0.51 = log ( [base] / [acid] )
[base] / [acid] = 100.51 = 3.24
The amount of acid or base a buffer solution can neutralize without a significant change in pH
Buffer capacity increases with increasing absolute concentration of the buffer solution (increasing concentration of components)
Buffer capacity also increases as the concentration of acid and conjugate base reach similar levels (the [base] / [acid] approaches one)
However, if a buffer neutralizes mainly acid or mainly base it may have a much higher concentration of one of the components
Buffer Capacity
Acid-Base titrations: an acidic (or basic) solution of unknown concentration reacts with a basic (or acidic) solution of known concentration in order to determine the original solution’s concentration
Titrations
Titrating
A pH indicator is mixed with the solution to monitor pHThe solution of known acid or base concentration is slowly added in using a buret The acid and base neutralize each other as the titration continuesWhen the equivalence point is reached the titration is complete
(pH indicator- dye whose color depends on pH)
(Equivalence point- point when the number of moles of acid and base are stoichiometrically equal)
Buret
Strong Acid-Strong Base Titration
Ex: Titration of 25 mL 0.100M HCl with 0.100M NaOHCalculating equivalence point:HCl (aq) + NaOH (aq) ⇒ H2O (l) + NaCl (aq)
Initial mol HCl = 0.025L x (0.100mol / 1 L) = 0.0025mol HClNaOH solution = 0.0025mol x (1L / 0.100mol) = 0.0250L
The pH at the equivalence point will always be 7 (all H3O
+ and OH- ions have neutralized each other)
To find the concentration of the initial solution convert the volume added to moles and then moles to concentration
MaVa = MbVb(mol acid) x (volume acid) = (mol base) x (volume acid)
( mol / L ) x (L) = ( mol / L ) x (L)
This equation is used to find the equivalence point
Strong Acid-Strong Base Titration ContinuedInitial pH= -log [H3O
+] = -log (0.100) = 1
pH at any given point before equivalence point:
mol NaOH added = (L added) x ( [NaOH] / 1 L )Because we are working with a strong acid and base H3O
+ = mol HCl - mol NaOH[H3O
+] = (mol H3O+ / total volume)
pH = -log [H3O+]
pH at any given point after equivalence point:
[OH-] = (mol OH- - mol H3O+) / total volume ⇒ [H3O
+] = 10-14 / [OH-] ⇒ pH = -log [H3O
+]
Practice
Calculate the pH after adding 5 mL of 0.100 M NaOH to 25 mL of 0.100 M HCl
0.00 0.0025 mol
0.0005 mol —
0.00 mol 0.0020 mol
Before Addition
Addition
After Addition
[H3O+] = ( .0020 mol H3O
+ ) / ( 0.0250 L + 0.00500 L ) = 0.0667
pH = -log 0.0667 = 1.18
OH- + H3O+ ⇒ 2 H2O
Convert mL ⇒ mol, Use addition table to find mol H3O+
mol H3O+ ⇒ [H3O
+], [H3O+] ⇒ pH
Total Volume
Titration curve/pH curve
Strong Base-Strong Acid Titration
College Board- Buffer
College Board- Titrations