BUFFERS
8.8
BUFFER
Half equivalence point:
pH = pKa
H+ = K a
• Weak acids and bases
always have a buffering
zone because of their
conjugate acid/base
Why doesn’t the acid
bas pair neutralize
eachother? (They
are conjugates)
Buffering zone: resists pH change when a
strong acid/base is added to a weak due
to conjugate acid-base pair
BUFFER
Buffering Capacity: is
the amount of H+ or
OH- the buffer can
absorb
• A buffer is an equal mix of a weak acid/base and its conjugate
– This is naturally made during a titration but due to lack of conjugate may have limited
capacity
• Example: H2CH3O2 and HCH3O2– A effective buffer can be made from a weak acid/base and adding a salt that contains
respective conjugate ion
• Example: HF and F- needs NaF(s) to be added to the solution
• Example: NH3 and NH4+ needs NH4Cl(s) to be added
We can make a buffer of any pH by controlling the amount of salt (conjugate ion) added
HOW DO BUFFERS WORK?
1. Add Base (OH-)
2. Add Acid (H+)
OH
H+- ++ HA
+ H+ + A - + H
OH
OH-+ HA
HA
A-
+
HH
O
+
A
+
2
This will push the reaction
towards the reactants
products thus
thus
“absorbing”
“absorbing”
somesome
of theof
the
baseacid
CALCULATION
Choose an acid with a
pKa close to the pH you
want to achieve
•
Henderson-Hasselbal
ch equation
EXAMPLE
• Hydrazine (N2H4) is a weak base with a Kb = 1.7 x 10-6. What compound could you add to a
0.50 mol/L solution to make a buffer solution?
• Extension: If the concentration of both the hydrazine and conjugate in the final solution were
0.5M what [H3O+] would exist in the buffer solution? What is the pH?
INVESTIGATION
• Go to the animation:
http://introchem.chem.okstate.edu/DCICLA/pHbuffer20.html
• Create a buffer solution using equal volumes of 0.10 M nitric acid and 0.10 M sodium nitrate.
• Insert the probes and record the pH (in your table)
• Remove the probes
• Click go to Part II of the simulation
• Add 0.001 moles of HCl - insert probes and record pH
• Add 0.001 moles of NaOH – insert probes and record pH
• Return to Part I and try the next combinations (in your table)
REAL WORLD EXAMPLES
• Blood plasma is a buffer. Your blood pH is 7.4, a change of 0.2 could be life threatening.
(Orange juice could kill you)
• Enzymes only function within a pH range in your body
• Fermentation and antibiotics require buffers to avoid side reactions.
EXAMPLE
• An environmental chemist requires a solution buffered to pH 9.8 to study the effects of acid
rain on aquatic microorganisms. Decide on the most appropriate butter components and
suggest their appropriate relative concentrations.