BIOL 1406 Austin Community College Understanding pH and Buffers Lab Report I NEED THIS TO BE DONE IN 7 HOURS NO LATER THAN THAT PLEASE !I need help with my

BIOL 1406 Austin Community College Understanding pH and Buffers Lab Report I NEED THIS TO BE DONE IN 7 HOURS NO LATER THAN THAT PLEASE !I need help with my BIO lab ,, I only need you to do the part that says Post lab only POST LAB. I couldn’t do it because I don’t know how to use Excel and on the post lab there is some parts that I will need to do it on excel Lab 3:
Understanding pH and Buffers
Objectives
After completing this exercise, you should be able to:
•
•
•
•
Define and correctly use the following terms: dissociation, acid, base, pH, buffer, buffering range, and
buffering capacity
+
Explain how solutions with different pHs compare with respect to their H and OH concentrations
Calibrate and use a pH meter to measure the pH of a solution
Plot and interpret a pH titration curve to determine the buffering range and buffering capacity of a
buffered solution
Prelab
Understanding pH and Buffers
A. Dissociation of Water
The chemistry of life is based largely on the chemistry of water. When ionic or polar solutes are mixed
with water, the polar water molecules will be attracted to them. Molecules held together by ionic bonds
may be pulled apart into oppositely charged ions. This process is called dissociation. In fact, water
+
molecules may even pull other water molecules apart to produce oppositely charged hydrogen ions (H )
and hydroxide ions (OH ):
H2O ↔ H+ + OHThe double-headed arrow in the equation above indicates that this is a reversible reaction and can occur
+
in both directions. Because each water molecule that dissociates produces one H and one OH , pure
+
-7
water always has the same concentration of H and OH , which turns out to be 1 × 10 M:
In pure water, [H+ ] = [OH- ] = 1 x 10-7 M
+
The square brackets means “concentration of”, so “[H ]” means “the concentration of hydrogen ions”.
+
Note that H goes by several names. It is often called a “hydrogen ion”, but it may also be called a
“hydronium ion” or a “proton”.
B. Acids and Bases
Although acids and bases have been defined in several different ways, in this class we will use the
Bronsted-Lowry definition of an acid as a proton donor, and a base as a proton acceptor. In the
following chemical equation, hydrochloric acid (HCl) is the acid and ammonia (NH3) is the base:
HCl + NH3
ACC BIOL 1406 Lab Manual Hays Campus Edition Fall 2018
Cl- + NH4+
Lab 3
Page 1
–
HCl dissociates and donates its lost proton to NH3, resulting in the formation of a chloride ion (Cl ) and an
+
ammonium ion (NH4 ). The reaction can go at a slower rate in the reverse direction, as indicated by the
+
smaller arrow pointing to the left. When this occurs, the NH4 becomes the proton donor and the Cl
becomes the proton acceptor.
+
Because acids are proton donors, they increase [H ] when they dissociate in water. For example, the
+
+
+
hydrochloric acid (HCl) dissociates to form H and Cl-. The released H increases the solution’s [H ] so
+
that it is higher than in pure water. As [H ] increases in a solution, the relative [OH ] will decrease.
–
+
Bases, on the other hand, decrease a solution’s [H ] in one of two ways:
–
•
Some bases dissociate to release OH . For example, sodium hydroxide (NaOH) is a strong base that
+
+
+
easily dissociates to form Na and OH . As [OH ] increases, [H ] decreases because some of the H
ions that were present before the NaOH was added will combine with the added OH- to form water.
•
Some bases combine with H ions, decreasing the amount of them in the solution. As the number of
+
free H decreases, more water molecules will dissociate and increase the number of OH ions. For
+
+
example, NH3 combines with H to form NH4 .
+
+
–
In summary, when an acid is added to an aqueous solution, [H ] increases and [OH ] decreases.
+
When a base is added to an aqueous solution, [OH ] increases and [H ] decreases.
Acids and bases can vary in their strength. Strong acids, such as hydrochloric acid (HCl) and sulfuric acid
+
(H2SO4), almost completely dissociate in water. Therefore, they cause a relatively large increase in [H ]
and a correspondingly large decrease in [OH ]. Weak acids, like acetic acid or citric acid, dissociate much
+
more slowly. This results in a relatively small increase in [H ] and a correspondingly small decrease in
[OH ].
Similarly, strong bases, such as sodium hydroxide (NaOH) and potassium hydroxide (KOH), produce a
large increase in [OH ] when dissolved in water, while weak bases such as ammonia (NH3) cause a much
smaller increase in [OH ]. Most biomolecules act as weak acids and/or bases.
C. pH
The pH of a solution is defined as the negative base-10 logarithm of the hydrogen ion concentration of
the solution:
pH = – log [H+ ]
NOTE: The pH will be a positive value. Since it is a logarithm, pH is also unitless.
In this equation, the base-10 logarithm (or “log”) is the power to which 10 must be raised to give the
desired number. For example, the log of 100 is 2 (to get 100, we must raise 10 to the power of 2) and the
log of 0.01 is -2 (to get 0.01, we must raise 10 to the power of -2).
+
Notice that the equation takes the negative log of the [H ]. This ensures that the pH will be a positive
-7
-7
value. A neutral solution has a hydrogen ion concentration of 1 × 10 M. The log of 1 × 10 is -7, so
-7
the negative log of 1 × 10 is 7. Therefore, the pH of a neutral solution is 7.
+
The [H ] of a solution is inversely related to its pH:
+
-7
+
-7
An acid has a [H ] greater than 1 x10 M and a pH lower than 7. A base has a [H ] less than 1 x10
M and a pH higher than 7.
Thus, a pH is more acidic the closer it is to zero and is more basic the closer it is to 14.
ACC BIOL 1406 Lab Manual Hays Campus Edition Fall 2018
Lab 3
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The Campbell Biology textbook’s representation of a pH scale is shown below:
To describe the relative strength of an acid or base, consider this example: Solution A has a pH of 3.
+
-3
+
-5
Thus, its [H ] = 0.001 M, or 1 x 10 M. Solution B has a pH of 5, so its [H ] = 0.00001 M, or 1 x 10 M.
+
Both solutions are acidic. Solution A is more acidic than B: Solution A has a lower pH and a greater [H ]
+
2
than Solution B. Furthermore, the [H ] of Solution A and B differ by 10 , or 100. Thus, you could say that
“Solution A is a 100 times stronger acid than Solution B” or “Solution A is 100 times more acidic than
Solution B”. You could also say instead that “Solution B is a 100 times stronger base than Solution A” or
“Solution B is 100 times more basic than Solution A”.
The difference between whole-number pH values is always 10 raised to an exponent!
+
If you are given the pH of a solution and are asked to calculate its [H ], you would basically reverse the
base-10 logarithm. This is sometimes called the antilogarithm, or “antilog”. An easy way to do this
without a calculator is to make the pH the negative exponent of 10. For example, a solution with a pH of
+
-10
+
-2
10 would have a [H ] of 1 x 10 M. A solution with a pH of 2 would have a [H ] of 1 x 10 M.
ACC BIOL 1406 Lab Manual Hays Campus Edition Fall 2018
Lab 3
Page 3
Your Turn
1. Determine the pH for each of the following solutions, given the [H+].
-3
[H+] = 1 x 10 M
-5.5
[H+] = 1 x 10
-8
M
[H+] = 1 x 10 M
-10
pH
[H+] = 1 x 10
M
pH
[H+] = 1 x 10 M
pH
[H+] = 1 x 10
-1
-13
M
pH
pH
pH
2. Determine the [H+] for each of the following solutions, given the pH.
pH 5
M
pH 2
M
pH 9
M
pH 12
M
pH 6
M
pH 7
M
3. Arrange the pH values given in Question 2 in order, from most acidic to most basic.
Most acidic
4.
then
then
then
then
then
Most basic
Complete the following:
a.
A solution of pH 5 has a
b.
A solution of pH 2 is
c.
A pH 9 solution is
times larger concentration of H+ than a solution of pH 6.
times more acidic than a solution of pH 8.
times more basic than a pH 7 solution.
D. Buffers
A buffer or buffer system is a mixture of molecules that maintain a certain pH (not necessarily pH 7) by
+
resisting pH changes. Buffers release or bind H to maintain a relatively constant pH. Most buffers consist
+
+
of a weak acid (which releases H ) and a weak base (which binds H .) If additional base is added
+
to a buffered solution, the weak acid of the buffer releases H , which combine with the extra OH to
form water. Therefore, the solution does not become as basic as it would have without the buffer. If
+
additional acid is added to a buffered solution, the weak base of the buffer combines with the H
+
released by the added acid. This reduces the concentration of free H in the solution so that it doesn’t
become as acidic as it would have without the buffer.
There are many different buffers, and each one will stabilize a solution so that it stays within a specific pH
range, called its buffering range. One buffer may be effective within a range of pH 2 to pH 6, while
another may be effective within a range of pH 10 to pH 12. Buffers are unable to stabilize a pH outside of
their buffering range.
Each buffer also has a certain buffering capacity, which is determined by the amount of additional acid
or base that the buffer can handle and still maintain its pH range. The buffering capacity is exceeded when
enough acid or base is added so that buffer is unable to maintain the desired pH range.
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Lab 3
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Buffers are extremely important to living organisms because most biochemical processes proceed
+
normally only when the pH remains within a fairly narrow range. An excess of H or OH can interfere with
the structure and function of many biomolecules, especially proteins. Therefore, buffers are commonly
used in living organisms to help maintain a relatively stable pH. In humans, buffers maintain blood pH
between 7.35 and 7.45 even though acids and bases are continually being added to and removed from
the blood as it travels through the body.
In the laboratory, molecular and cellular biologists make extensive use of buffers to mimic in a test tube
the buffered conditions of living cells and tissues. This ensures that the experiment is as close as
possible to the living, or in vivo, system.
Watch this video: Acids, Bases, & Buffers
E. pH Titration Curves
A pH titration curve is a graph that shows the pH change of a solution as increasing amounts of acid or
base are added. The X-axis of the graph shows the volume of acid or base added. Note that the 0 mL
mark is somewhere in the middle of the axis. This corresponds to the starting pH of the solution, before it
is titrated. The Y-axis shows the pH of the solution. You should be able to determine the buffering range
and the buffering capacity of a solution from a titration curve.
Buffering range is the pH range in which the buffer is effective at resisting pH changes, even if
additional acid or a base is added. The buffering range is the pH range where the titration curve appears
most horizontally flat. When the buffering range is exceeded, the solution can no longer resist the pH
changes caused by the increasing amount of acid or base and its pH drops or rises dramatically. Some
solutions can have more than one buffering range, as shown by the titration curve for Solution B.
Buffering capacity is the ability of a buffer to resist changes in pH when additional acid or base is added.
The capacity is determined by the relative length of the horizontal (flat) portion of the titration curve. For
this class, a solution’s buffering capacity is a qualitative measurement – it can be described as
relatively small or large, instead of assigning a specific value.
Two examples of titration curves are shown below. The data points have been left out for simplicity:
Titration curve for Solution A
(One buffering range)
Titration curve for Solution B
(Two buffering ranges)
14
14
13
13
12
12
11
11
10
10
9
9
8
8
pH 7
pH 7
6
6
5
5
4
4
3
3
2
2
1
1
5 4 3 2 1 0 1 2 3 4 5
HCl added (mL)
NaOH added (mL)
ACC BIOL 1406 Lab Manual Hays Campus Edition Fall 2018
5 4 3 2 1 0 1 2 3
HCl added (mL)
NaOH added (mL)
Lab 3
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Your Turn
Titration curve for Solution A
Titration curve for Solution B
14
14
13
13
12
12
11
11
10
10
9
9
8
8
pH 7
pH 7
6
6
5
5
4
4
3
3
2
2
1
1
5 4 3 2 1 0 1 2 3 4 5
HCl added (mL)
NaOH added (mL)
5 4 3 2 1 0 1 2 3
HCl added (mL)
NaOH added (mL)
Compare the titration curves for Solution A and Solution B.
What is the buffering range of Solution A?
What are the buffering ranges of Solution B?
Which Solution has a larger buffering capacity: A or B?
F. Making Buffered Solutions
A buffered solution is any solution that also contains the acid and base components to make up a buffer.
Without the acid and base components, the solution would not be able to maintain pH very well or at all.
Buffer solutions are important in biology labs and can be complex mixtures. The buffered solution that is
used to keep cells alive in a laboratory culture contains up to thirty nutrients, including vitamins and
amino acids, in addition to the acid and base that make up the buffer.
In this lab, you will make three solutions that all have the same concentration of sucrose. One will be an
unbuffered sucrose solution. The other two sucrose solutions will also contain the acid and base
components for an acetate buffer and a bicarbonate buffer respectively. You will be provided with stock
solutions to make all three of them. After you make these solutions, you will add acid and base to them to
observe how the buffers work.
ACC BIOL 1406 Lab Manual Hays Campus Edition Fall 2018
Lab 3
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Your Turn
In this box, you will calculate the volumes you will need to make the three sucrose solutions. The
following stock solutions will be provided to you in the lab:
0.1 M sucrose
0.2 M acetic acid
0.2 M sodium acetate
0.2 M anhydrous sodium carbonate
0.2 M sodium bicarbonate
Use the parallel dilution equation C1V1 = C2V2 to calculate V1 of each solute for the three sucrose
solutions described below.
For solutions that have more than one solute, add together the V1 you calculate for each one, and then
subtract this sum from the total desired volume to calculate the volume of dH2O that would be needed.
Don’t forget to include the units for all of your answers!
The first solution will be an unbuffered sucrose solution. Calculate the amount of 0.1 M sucrose
stock solution and the amount of dH2O needed to prepare 100 mL of unbuffered 20 mM sucrose.
Amount of 0.1 M sucrose stock needed
Amount of dH2O needed to bring to final volume
The second sucrose solution will contain sucrose, as well as acetic acid and sodium acetate to
create an acetate buffer of pH approximately 4.8. Calculate the amount of each stock solution and the
amount of dH2O needed to prepare 100 mL of a solution that contains 20 mM sucrose, 40mM acetic
acid, and 60 mM sodium acetate.
Amount of 0.1 M sucrose stock needed
Amount of 0.2 M acetic acid stock needed
Amount of 0.2 M sodium acetate stock needed
Amount of dH2O needed to bring to final volume
The third sucrose solution will contain sucrose, as well as sodium carbonate and sodium
bicarbonate to create a bicarbonate buffer of pH approximately 9.8. Calculate the amount of each
stock solution and the amount of dH2O needed to prepare 100 mL of a solution that contains 20
mM sucrose, 22 mM anhydrous sodium carbonate, and 28 mM sodium bicarbonate.
Amount of 0.1 M sucrose stock needed
Amount of 0.2 M anhydrous sodium carbonate stock needed
Amount of 0.2 M sodium bicarbonate stock needed
Amount of dH2O needed to bring to final volume
Transfer all of your answers to Part I of the Lab Procedures.
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Lab 3
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Lab Procedures
Watch this video: Procedures for pH and buffers portion
I. Prepare Unbuffered and Buffered Solutions of Sucrose
A. Making 100 mL of Unbuffered Sucrose
Transfer the amounts you calculated as part of the Prelab to the blanks below. Don’t forget the units!
Amount of 0.1 M sucrose stock needed
Amount of dH2O needed to bring to final volume
Measure each solution, choosing the most appropriate measuring devices to accurately make the
solution in a 150 mL beaker labeled “20 mM unbuffered sucrose” with labeling tape. Label any glass
pipettes used with colored tape.
Before transferring half of the 20mM unbuffered sucrose solution, measure the pH of this solution and
add these values to Table 1 and Table 4 for the 0 mL value.
B. Making 100 mL Acetate Buffer with Sucrose
Transfer the amounts you calculated as part of the Prelab to the blanks below. Don’t forget the
units!
Amount of 0.1 M sucrose stock needed
Amount of 0.2 M acetic acid stock needed
Amount of 0.2 M sodium acetate stock needed
Amount of dH2O needed to bring to final volume
Measure each solution, choosing the most appropriate measuring devices to acc urately make the
solution in a 150 mL beaker labeled “20 mM acetate buffered sucrose” with labeling tape. Label any
glass pipettes used with colored tape.
Before transferring half of the 20mM acetate buffered sucrose solution, measure the pH of this
solution and add these values to Table 2 and Table 5 for the 0 mL value.
C. Making 100 mL Bicarbonate Buffer
Transfer the amounts you calculated as part of the Prelab to the blanks below. Don’t forget the units!
Amount of 0.1 M sucrose stock needed
Amount of 0.2 M anhydrous sodium carbonate stock needed
Amount of 0.2 M sodium bicarbonate stock needed
Amount of dH2O needed to bring to final volume
Measure each solution, choosing the most appropriate measuring devices to accurately make the
solution in a 150 mL beaker labeled “20 mM bicarbonate buffered sucrose” with labeling tape. Label
any glass pipettes used with colored tape.
ACC BIOL 1406 Lab Manual Hays Campus Edition Fall 2018
Lab 3
Page 8
Before transferring half of the 20mM bicarbonate buffered sucrose solution, measure the pH of this
solution and add these values to Table 3 and Table 6 for the 0 mL value.
II. Titrating Sucrose Solutions with HCl (Strong Acid)
A. Calibrate the pH Meter
1.
2.
Before the meter can be used, it must be calibrated. This only has to be done once during the lab.
Calibrate the pH meter while the electrode is sitting in the green pH 7 standard buffer:
a.
The pH probe has been set and calibrated to a pH of approximately 7.00 for you by lab tech.
b.
If the value does not read approximately 7.00, press the red box and select “Zero” from the drop
down menu.
c.
The device should now read approximately 7.00
Some important notes about using the pH meter:
IMPORTANT: To avoid cross-contamination, always rinse the electrode (measuring part of the
pH meter) with dH 2O from a wash bottle before transferring it to another solution. Hold the
electrode over a waste beaker while you rinse.
IMPORTANT: The measuring end of the electrode should always be sitting in liquid. The
electrode might be permanently damaged if you allow it to dry out. Notice that it is currently
sitting in a green pH 7 standard buffer. If you are not using the electrode to take measurements,
always rinse it with dH2O and return to this buffer.
B. Titrating Unbuffered Sucrose
1.
Transfer half (50 mL) of the unbuffered 20 mM sucrose solution to a 100 mL beaker. Save the
remaining solution for Part III of the experiment (Titrating Sucrose Solutions with NaOH on pg.
13). Place the 100 mL beaker labeled “20 mM unbuffered sucrose” on the stirrer and add a small
magnetic stir bar. Turn on the stirrer and gradually adjust the rate so that it is a moderate speed. The
solution should not be splashing.
2.
Remove the electrode from the standard buffer and rinse it with dH2O from a wash bottle. Hold the
electrode over the waste rinse water beaker while doing this.
3.
Slowly lower the electrode into the unbuffered 20 mM sucrose solution. Place the pH meter in the
beaker so that it stably tilts against the edge of the beaker and not being hit by the spinning stir bar.
4.
Wait about 10 seconds or more until the reading stabilizes and record the pH value that appears
on the display. This is the pH of the solution when 0 mL of HCl has been added. Record pH in Table
1.
5.
Using micropipettor, add 1 mL of 0.4 M HCl to the sucrose solution. D o not let t he H Cl t ou…
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