Question

A) Need to prepare 100 ml of a 0.10 M phosphate buffer with a pH of 6.2. Calculate the amounts of 1.0 M NaH2PO4 and 1.0 M K2HPO4 solutions which will give the required pH of 6.2.Use the Henderson-Hasselbalch equation ( pH = pKa + log (base/acid)

B) Need to prepare 100 ml of a solution containing 0.025 M NaHCO3 and 0.12 M NaHCO3 and 0.12 M NaCl From stock solutions of 0.05 M NaHCO3 and 0.24 M NaCl from stock solutions of 0.05 M NaHCO3 and 0.24 M NaCl. Calculate how much of each solution you will need.

Answer #1

a) The pKa of the reaction
H_{2}PO_{4}^{-} ? H^{+} +
HPO_{4}^{2-} is 7.21 (source wikipedia
http://en.wikipedia.org/wiki/Phosphoric_acid)

Thus using the Henderson-Hasselbalch equation we arrive at
[HPO_{4}^{2-}]/[
H_{2}PO_{4}^{-} ] = 0.1

Therefore, 0.1 M phosphate buffer at pH 6.2 will have 9.1%
of H_{2}PO_{4}^{-} and 91% of
HPO_{4}^{2-} OR in 100ml of buffer 0.0091 mol and
0.091 mol respectively.

Thus mixing 91 ml of 1.0 M K_{2}HPO_{4} and 9.1
ml of 1.0 M NaH_{2}PO_{4} solution results in 100
ml of 0.1 M phosphate buffer

b) I think the question in correct form is 100 ml of a solution
containing 0.025M NaHCO_{3} and 0.12 M NaCl from stock
solutions of 0.05 M NaHCO_{3} and 0.24 M NaCl.

Since the required solution contains the concentration half of
the stock solution, mixing of 50 ml each from the stock solution
will result 100 ml of solution containing 0.025M NaHCO_{3}
and ).12M NaCl.

You need to prepare a buffer for biochemistry lab. The required
solution is 0.5Msodium phosphate, pH 7.0. Use the
Henderson-Hasselbalch equation to calculate the number of moles and
grams of moobasic sodium phosphate ( NaH2PO4) and dibasic sodium
phosphate ( Na2HPO4) necessary to make 1 liter of solution. The pKa
for this buffer is 7.21

1. Show the calculations required for preparing the buffers in
this exercise. The pKa values for H3PO4 are 2.15, 6.82, 12.32. Note
that you’ll be using the pKa value of 6.82. Why?
Method 1 a. Calculate the grams of NaH2PO4*H2O needed to prepare
100.0 mL of a 100mM solution. b. Calculate the amount of 1.00M NaOH
you expect to add to adjust the pH of your NaH2PO4 solution to
7.0.
Method 2 c. Using the Henderson-Hasselbalch equation, calculate
the g...

1) You need to prepare 1.000L of 0.50 M phosphate
buffer, pH 6.21. Use the Henderson-Hasselbalch equation with a
value of 6.64 for pK2, to calculate the quantities of K2PO4 and
KH2PO4 you need to add to your flask.
What is the ratio you need of [K2HPO4]/[KH2PO4] .
For this question I know there is an answer already up, but it
doesn't show work to help me understand how they got that answer. I
also got something different using the...

1. You need to prepare 1.000 L of 0.50 M phosphate buffer,
pH=6.21. Use the Henderson-Hasselbalch equation with a value of
6.64 for pK2 to calculate the quantities of K2HPO4 and K2H2PO4 you
need to add to the flask. Record the steps of these calculations
and use them as a guide for the calculation you will have to make
in the laboratory. (Calculate the intermediate values to at least
one more significant figure than required.)
What is the [K2HPO4]/[K2H2PO4] ratio...

1. You need to prepare 1.000 L of 0.50 M phosphate buffer,
pH=6.21. Use the Henderson-Hasselbalch equation with a value of
6.64 for pK2 to calculate the quantities of K2HPO4 and K2H2PO4 you
need to add to the flask. Record the steps of these calculations
and use them as a guide for the calculation you will have to make
in the laboratory. (Calculate the intermediate values to at least
one more significant figure than required.)
What is the [K2HPO4]/[K2H2PO4] ratio...

The Henderson-Hasselbalch equation relates the pH of a buffer
solution to the pKa of its conjugate acid and the ratio of
the concentrations of the conjugate base and acid. The equation is
important in laboratory work that makes use of buffered solutions,
in industrial processes where pH needs to be controlled, and in
medicine, where understanding the Henderson-Hasselbalch equation is
critical for the control of blood pH.
Part A
As a technician in a large pharmaceutical research firm, you need...

0.10 M Phospate buffer at pH 7.0 total 100.0 ml. 30mL of 0.20M
NaH2PO4, 20 mL of 0.20M Na2HPO4 and 50.0 mL of water were used to
prepare the buffer.
The average pKa is 6.76
Using the average pKa value calculated in 3, calculate the pH of
the pH 7.0 buffer solution that is expected after the addition of
1.0 mL of 1.0M HCl.

To prepare 100 mL of 25 mM Tris buffer at pH 8.0, calculate
the amount of Tris base and Tris·HCl needed to make the buffer.
(Hint: find out the pKa of Tris·HCl and use the
Henderson-Hasselbalch Equation.) Describe how do you want to make
the solution, including how much water to add, how to adjust pH,
etc.

You wish to prepare 100 mL (total volume) of a buffer solution
that is 0.025 M in carbonic acid, pH 7.25. You have solid H2CO3 and
solid NaHCO3. Calculate how many grams of H2CO3 and HCO3- you need
to weigh out to prepare this solution.
This question has been posted before but can someone please
explain where the 7.76, 8.76, and 1 came from?
7.25=6.36 +log (base/acid)
0.89= log (base/acid)
7.76=base/acid
base=(7.76/8.76) *0.025M/0.1L=0.2214M NaHCO3
acid=(1/8.76) * 0.025M/0.1L=0.0285M H2CO3
Need 0.1L*0.025M...

The Henderson-Hasselbalch equation relates the pH of a buffer
solution to the pKa of its conjugate acid and the ratio of
the concentrations of the conjugate base and acid. The equation is
important in laboratory work that makes use of buffered solutions,
in industrial processes where pH needs to be controlled, and in
medicine, where understanding the Henderson-Hasselbalch equation is
critical for the control of blood pH.
Part A.) As a technician in a large
pharmaceutical research firm, you need...

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