LAB 5. ACIDS AND BASES:
PH AND BUFFERS
PURPOSE: To determine the pH of common acids and bases using a pH meter, pH paper, and
red cabbage indicator.
To test the effect of adding an acid or base to a buffer solution.
SAFETY CONCERNS:
Always wear safety goggles.
Wash with soap and water if skin contacts acids or bases.
ACIDS:
An Acid is a substance that when dissolved in water will produce hydrogen ions, H+, in the
solution. An acid that does not contain carbon is called an inorganic acid. An acid that contains
carbon is called an organic acid.
Strong acids are acids that produce lots of hydrogen ions, H1+’s. They almost completely
dissociate (break apart) to form hydrogen cations, H1+, and the companion anions. Once
dissolved in water the very strong acid does not exist any more since it has dissociated into its
ions.
H-Cl  H1+ + Cl1-
Dissociation of Hydrochloric acid
Examples of Strong Acids:
Strong Acid
Hydrochloric Acid
Sulfuric Acid
Formula
HCl
Common Source
H2SO4
Battery Acid
Stomach Acid
Weak acids are acids that produce only a few hydrogen ions, H1+’s. Only some of the
molecules in solution dissociate (break apart) into hydrogen cations, H1+, at a time. When they
do break apart, the hydrogen cations, H1+, and the companion anions join back together again to
reform the acid. This process of the acid breaking apart and then reforming over and over is
called equilibrium. Because of the equilibrium there is a mixture of the original acid, the
hydrogen ion, H1+, and the anion all in the solution at the same time.
Dissociation of Acetic acid
H-C2H3O2 H1+ + C2H3O21-
Dissociation of
Carbonic acid
CO2 + H2O  H2CO3 H1+ + HCO31CH110 Lab 5. Acids and Bases (W14)
51
Examples of Weak Acids:
Weak Acid
Acetic Acid
Formula
HC2H3O2
Common Source
Carbonic Acid
H2CO3
In Carbonated Water
Citric Acid
H3C6H5O7
In Lemons & Oranges
Tartaric Acid
H2C4H4O6
In Grapes
Phosphoric Acid
H3PO4
In Cola Drinks
Vinegar
(A weak inorganic acid)
BASES:
A Base is a substance that when dissolved in water will produce hydroxide ions, OH1-, in the
solution. A base that does not contain carbon is called an inorganic base. A base that contains
carbon is called an organic base.
Strong bases are bases that produce lots of hydroxide ions, OH1-. They almost completely
dissociate (break apart) to form hydroxide anions, OH1-, and the companion cations.
Dissociation of
Sodium Hydroxide
Examples of Strong Bases:
Strong Base
Sodium Hydroxide
NaOH  Na1+ + OH1-
Formula
NaOH
(strong)
Potassium Hydroxide
KOH
Common Source
Lye, Caustic Soda,
Soda Ash, Drano
Potash
(strong)
Weak bases are bases that produce only a few hydroxide ions, OH1-. Sometimes it’s not
obvious that hydroxide anions, OH1-, are even produced, but the anions that are produced can
react with water to produce OH1-’s. Weak bases are in equilibrium with their ions.
Dissociation of
Sodium
Bicarbonate
52
NaHCO3  Na1+ + HCO31-

HCO31- + H2O  H2CO3 + OH1-
CH110 Lab 5. Acids and Bases (W14)
Examples of Weak Bases:
Weak Base
Magnesium Hydroxide
Formula
Mg(OH)2
Common Source
Ammonium Hydroxide
NH4OH
In Glass Cleaner
Sodium Bicarbonate
NaHCO3
Baking Soda
Calcium Carbonate
CaCO3
Antacids, Sea Shells,
Egg Shells,
Limestone & Marble
In Milk of Magnesia
PH:
The pH of a solution is an indicator of the number of hydrogen ions, H1+’s, present in a solution.
We measure the concentration of hydrogen ions, H1+’s, in moles per liter, M, which we
symbolize with square brackets, [H1+].
Pure water contains a small amount of both hydrogen ions, and hydroxide ions. In pure water
the concentrations of hydrogen ion and hydroxide ions are the same.
[H1+] = [OH1-] = 0.000,000,1 M.
It is awkward to work with numbers as small as 0.000,000,1 M. It helps to report the hydrogen
ion concentration, [H1+], in scientific notation. To simplify reporting the concentration of H1+’s
even more, we use the positive value of the exponent only and call it the pH.
(in scientific notation)
1+
1+
[H ]
0.000,1 M
0.000,000,1 M
0.000,000,000,1 M
(log means opposite of exponent on base 10)
-pH
pH = -log [H1+]
4
7
10
[H ] = 1x10
1 x 10-4 M
1 x 10-7 M
1 x 10-10 M
If the [H1+] = [OH1-] then the solution is considered to be neutral and the pH = 7
If the [H1+] > [OH1-] then the solution is considered to be acidic and the pH < 7
If the [H1+] < [OH1-] then the solution is considered to be basic and the pH > 7
[H1+]= 10-1
pH =
1

10-2
10-3
10-4
10-5
10-6
2
3
4
5
6
Increasingly Acidic
[H1+] > [OH1-]
10-7
10-8
10-9
10-10
10-11
10-12
10-13
10-14
7
8

Neutral
9
10
11
12
13
14
[H1+] = [OH1-]
CH110 Lab 5. Acids and Bases (W14)
Increasingly Basic 
[H1+] < [OH1-]
53
The [H1+] will not always be a simple 1 x 10whole number value. In such cases we can estimate the
pH range or we can calculate the pH exactly using the pH equation: pH = -log [H+].
Examples:
[H1+] = 1x10-pH
pH = -log [H1+]
Given [H1+]
If [H1+] = 0.000,032M = 3.2 x 10-5M
Calculated [H1+]
Then [H1+] = 1 x 10-8.6 = 2.5 x 10-9 M
We can
estimate
pH to be
between 4 & 5
We can
estimate [H1+]
to be between
10-8 & 10-9
Calculated pH
Then pH = -log(3.2x10-5) =
4.5
Given pH
If pH = 8.6
INDICATORS:
The pH of a solution is often measured by observing how the acid or base causes the color of
certain organic molecules to change.
Litmus, a chemical found in lichen, is one of many acid-base indicators. Litmus turns
from blue to red in acidic solutions and from red to blue in basic solutions.
Phenolphthalein, a laxative, is colorless when acidic but turns brilliant pink above pH 7.
Anthocyanin indicators are common in flower petals, berries, and purple cabbage. The
colors change over the entire spectrum of pH making it a universal indicator. The chart
below shows the relative color changes of anthocyanin pigments over a wide pH range.
Colors of Anthocyanin (Red Cabbage) Indicator
pH = 1
2
Color Red
3
4
5
Fushia
Rose
6
7
Purple
(neutral)
8
9
10
11
12
13
14
Blue
Green GreenYellow
Aqua
brown
NEUTRALIZATIONS:
A Neutralization reaction is a reaction between a strong acid and a strong base in which the
product formed is neither acidic nor basic. For example, if hydrochloric acid (an acid) is mixed
with sodium hydroxide (a base), the products will be sodium chloride (table salt) and water.
Neutralization:
Strong Acid + Strong Base  Salt + Water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O

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CH110 Lab 5. Acids and Bases (W14)
BUFFERS:
A Buffer is a solution that will not drastically change in pH even when strong acids or bases are
added to it. A buffer is made by combining a weak acid and a weak base in a water
solution.
By having both an acid and a base in the solution together there will always be an acid to react
with any base added, or there will be a base to react with any acid added. For example, a simple
buffer could be made by combining equal amounts of the weak acid, Acetic Acid, and the weak
base, Sodium Acetate.
Buffer = Weak Acid + Weak Base
HC2H3O2(aq)
+
NaC2H3O2(aq)
(weak acid)
(weak base)
+
NaOH
+
HCl
If a strong base gets added to the buffer, it reacts with
the weak acid present
(rather than giving the solution more OH-‘s)
If a strong acid gets added to the buffer, it reacts with
the weak base present
(rather than giving the solution more H+’s)
and gets turned into
and gets turned into
NaC2H3O2(aq)
HC2H3O2(aq)
a weak base that does not drastically change the
pH of the solution .
a weak acid that does not drastically change the
pH of the solution.
As long as the number of OH1-‘s and H1+’s does not drastically change then the pH does not
drastically change. Buffer solutions are important in keeping the pH of biological solutions in a
narrow range.
CH110 Lab 5. Acids and Bases (W14)
55
PROCEDURES:
ACTIONS:
NOTES:
I. DETERMINATION OF PH OF COMMON ITEMS:
1. Arrange your large test tubes and into each one put a different
sample of a liquid1 (about 1 inch deep) or a solid (about a dime
sized scoop) of household product to be tested.
2. Add to each sample enough deionized water so that the probe
of your pH meter can just be immersed into the liquid. Stir or
stopper and shake each tube to dissolve.
3. Use these sample tubes for the sections A, B, & C where the
pH of each will be determined by various methods.
A. By Meter:
1. Clean the pH electrodes of your pH meter with deionized
water between each sample test.2 Be gentle with the electrodes
as they are delicate and can easily break or scratch.
2. Carefully submerge the cleaned electrode of a pH meter into
the samples indicated. 3 Read and record the pH on your report
sheet to the accuracy of your meter.
3. Calculate the [H1+] from your pH meter reading and record it
on your report sheet.
B. By Paper:
1. Tear the universal indicator pH paper into 1 cm square
pieces and arrange them on a watch glass or paper towel.4
2. Dip a glass stirring rod into the sample to be tested and touch
the wet stirring rod to the pH paper.
3. Compare the color produced on the wet5 paper to the color on
the pH paper container label. Record the results.
C. By Cabbage (Anthocyanins):
1. Tear a piece of purple cabbage6 into small pieces and place in
a glass beaker. 7
2. Add deionized water to barely cover the cabbage and then
boil it using a hot plate, burner, or microwave oven until the
water is dark purple. Remove from heat.
3. Into the test tube or beaker of the sample to be tested, 8 add a
few mLs of purple cabbage juice and observe the color change.
Record the pH range. 9
56
CH110 Lab 5. Acids and Bases (W14)
1
If the liquid is highly viscous,
(thick like shampoo) treat it as
a solid.
2
We
must
have
clean
electrodes for each test. If you
do not rinse off the previous
sample then the new sample
gets contaminated and the pH
can be altered.
3
Toilet cleaner and drano/lye
are very strong and so tax the
flexibility of our meters. To
avoid recalibrating our meters
we will test these with other
methods.
4
Bring the sample to be tested
to the pH paper rather than
inserting the pH paper into the
solution. Putting the pH paper
into the solution washes
pigment from the paper into
the sample.
5
Compare the color when the
paper is wet. The color will
change as the paper dries and
not compare accurately.
6
Roses,
hydrangeas,
and
blueberries
also
contain
anthocyanin pigments and will
work as indicators. We use red
cabbage because it is so
inexpensive
and
readily
available.
7
If working at home use a
coffee mug or glass container.
Metal pans cause the purple
pigments to turn blue.
8
Because tea and coffee are so
dark in color it is difficult to
see any change in the cabbage
pigment. We will rely on other
methods to test these.
9
Compare the colors of your
substance-cabbage mix with
the colors on the chart in the
discussion section. This is not
exact but just allows you to
predict a general pH range.
II. PH OF CARBON DIOXIDE, CO2, SOLUTION:
1. Place about 10 mLs of deionized water into a large test
tube or small beaker and determine the pH using a pH
meter.10
2. Keeping the pH electrode in the tube, place a straw in
the water and blow bubbles into it.11 Watch the pH
meter to observe changes.
3. Continue to blow bubbles into the water for a couple of
minutes. Record any pH changes on the report sheet.
III. BUFFERS:
10
The deionized water may not be
at a pH of 7 depending on the
amount of CO2 gases already
dissolved in it. Record the pH as
you find it and use this as your
starting point.
11
As you blow into the tube you
are adding Carbon Dioxide, CO2
to the water. CO2 + H2O forms
H2CO3. Review the disassociation
of Carbonic acid, H2CO3, in the
laboratory discussion of weak
acids.
1. Obtain 4 large test tubes and label them Tubes A-D.
Into both tubes A and B place about 10 mLs of
deionized or distilled water.
Into both tubes C and D make buffer solutions by
combining
5 mLs of 0.1 M Sodium Acetate, (NaC2H3O2 ) with
5 mLs of 0.1 M 0.1 M Acetic acid, (HC2H3O2) and
mix well.
A. Water with Acid:
2. Measure and record the pH of the water in tube A using
a meter. 10
3. Add 1 drop of 1.0 M HCl, Hydrochloric Acid, and swirl
the sample to mix. Record the pH.
4. Continue to mix in 1 drop of HCl at a time, recording the
pH after each drop until a total of 5 drops have been
added. Discard the solution in the sink.
B. Water with Base:
5. Measure and record the pH of the water using a meter.10
6. Add 1 drop of 1.0 M NaOH, Sodium Hydroxide, and
swirl the sample to mix. Record the pH.
7. Continue to mix in 1 drop of NaOH at a time, recording
the pH after each drop until a total of 5 drops have been
added. Discard the solution in the sink.
8. Analyze the results and report your conclusions on the
report sheet.
CH110 Lab 5. Acids and Bases (W14)
57
12
C. Buffer with Acid:
9. Measure and record the pH of the buffer solution in
tube C using a meter. 12
10. Add 1 drop of 1.0 M HCl, Hydrochloric Acid, and swirl
the sample to mix. Record the pH.
11. Continue to mix in 1 drop of HCl at a time, mixing and
recording the pH after each drop until a total of 5 drops
have been added. Discard the solution in the sink.
D. Buffer with Base:
12. Measure and record the pH of the buffer solution in
tube D using a meter. 12
13. Add 1 drop of 1.0 M NaOH, Sodium Hydroxide, and
swirl the sample to mix. Record the pH.
14. Continue to mix in 1 drop of NaOH at a time, recording
the pH after each drop until a total of 5 drops have been
added. Discard the solution in the sink.
15. Analyze the results and report your conclusions on the
report sheet.
58
CH110 Lab 5. Acids and Bases (W14)
Not all buffers are neutral. The
pH depends on the weak acid and
the weak base used. Each buffer
will have its own unique pH.
Record the pH of yours and use this
as your starting value.
LAB 5. ACIDS AND BASES:
PRE LAB EXERCISES:
NAME_____________
DATE______________
_____1.
A solution has a pH of 8.7. The solution is ______
A. Acidic
B. Basic
C. Neutral
_____2.
Which solution has the smaller hydrogen ion concentration?
A. Solution A which has a hydrogen ion, [H1+], concentration of 3.8 x 10-8.
B. Solution B which has a hydrogen ion, [H1+], concentration of 2.3 x 10-4.
C. Not enough information to predict.
_____3.
Which solution has the lower pH?
A. Solution A which has a hydrogen ion, [H1+], concentration of 3.8 x 10-8.
B. Solution B which has a hydrogen ion, [H1+], concentration of 2.3 x 10-4.
C. Not enough information to predict.
_____4.
Which solution has the higher pH?
A. 0.1 M HCl.
B. 0.1 M HC2H3O2
C. Neither as both solutions have the same pH.
_____5.
Sodium hydroxide turns purple cabbage to a yellow color. Ammonium hydroxide
turns purple cabbage green. Which solution has the lower pH?
A. 0.1 M NaOH.
B. 0.1 M NH4OH
C. Neither as both solutions have the same pH.
_____6.
A buffer is ______
A. a solution of acid and base that should always have a pH of 7.
B. formed when a strong acid neutralizes a strong base.
C. a solution of a weak acid and a weak base that resists change in pH.
D. all of these.
7. Write the equation of Hydrochloric acid (HCl) with sodium acetate (NaC2H3O2) in a buffer
solution.
8. Write the equation of Sodium Hydroxide (NaOH) with acetic acid (HC2H3O2) in a buffer
solution.
CH110 Lab 5. Acids and Bases (W14)
59
60
CH110 Lab 5. Acids and Bases (W14)
LAB 5. ACIDS AND BASES:
NAME___________________
PARTNER_________DATE___
REPORT:
I: Determination of pH by Meter, Paper, and Cabbage:
Solution
A. pH by
Meter
Water (for comparison):
[H1+] = 10-pH
B. pH by
Paper
C. New Cabbage
pH by
color produced Cabbage
Deionized water used
Foods:
Soda Pop (light color)
________brand
Tea or Coffee
OMIT as color is too dark to see a
difference with cabbage
(circle which)
Orange Juice
Vinegar
Cleaners:
Shampoo
Detergent
Ammonia
Bleach
Toilet Cleaner
Do not insert the meter into this solution
Lye
Do not insert the meter into this solution
Others:
Baking Soda
(NaHCO3)
Milk of Magnesia
(MgOH)
Antacid
_______brand
Results Summary:
1.___ In general, foods and beverages are mostly _____________
A. Acidic
B. Basic
C. Neutral
D. there is no general trend
2.___ In general, cleaning supplies are mostly _____________
A. Acidic
B. Basic
C. Neutral
D. there is no general trend
Conclusion/Explanation/Analysis: Why were the results as they were?
3.
Explain any anomalies.
Show the reaction with stomach acid that explains how a Tums antacid (CaCO 3) settles an upset stomach
and gives off gas that make you burp.
CH110 Lab 5. Acids and Bases (W14)
61
II. pH of Carbon Dioxide, CO2, in water:
pH of Deionized Water = __________
pH of “Breath Carbonated” Water = ___________
Complete and balance the equation for the reaction of carbon dioxide with water to form carbonic acid.
___CO2 +
___H2O

Conclusion/Explanation of Results: (Why did the CO2 from your breath have the effect that it did?)
Results Summary:
1.___
The pH of deionized water may not be exactly 7.00 because _____ (Based on what should have occurred above.)
A. there is dissolved carbon dioxide gas in the water that forms carbonic acid.
B. there are dissolved hard water ion impurities in the water.
C. water is naturally acidic due to the 1x10 -7 M of H1+ ions present.
D. all of these.
2.___
Which of the following is true about the solubility of substances in water? (Based on what should have occurred
above and the principles of gas solubility.)
A.
B.
C.
D.
Increasing the temperature has no effect on the quantity of a gas that will dissolve in water.
Increasing the temperature will usually decrease the quantity of a gas that will dissolve in water.
Increasing the temperature will usually increase the quantity of a gas that will dissolve in water.
There is no general trend regarding temperature and gas solubility.
3.
A sample of deionized water had a pH of 6.0 until it was boiled. Once this water was boiled and cooled it
had a pH of 7? Explain how this could happen. (Based on principles of gas solubility.)
4.
Bacteria in the mouth converts sugar into acids that are capable of dissolving tooth enamel. Repeated
drinking of carbonated beverages, even those that do not contain sugar, can also contribute to loss of
enamel. Explain why carbonated beverages can act in this way and cause dental caries (tooth decay).
62
CH110 Lab 5. Acids and Bases (W14)
III. Buffers:
pH
Solution:
A. Water with Acid
pH
Solution:
C. Buffer with Acid
Deionized Water
After 1 drop 1.0M HCl
After 2 drops 1.0M HCl
After 3 drops 1.0M HCl
After 4 drops 1.0M HCl
After 5 drops 1.0M HCl
NaC2H3O2:HC2H3O2 Buffer
After 1 drop 1.0M HCl
After 2 drops 1.0M HCl
After 3 drops 1.0M HCl
After 4 drops 1.0M HCl
After 5 drops 1.0M HCl
B. Water with Base
D. Buffer with Base
Deionized Water
After 1 drop 1.0M NaOH
After 2 drops 1.0M NaOH
After 3 drops 1.0M NaOH
After 4 drops 1.0M NaOH
After 5 drops 1.0M NaOH
Conclusion/Explanation/Analysis: (Why were the
NaC2H3O2:HC2H3O2 Buffer
After 1 drop 1.0M NaOH
After 2 drops 1.0M NaOH
After 3 drops 1.0M NaOH
After 4 drops 1.0M NaOH
After 5 drops 1.0M NaOH
Conclusion/Explanation/Analysis: (How does the
results as they were? When adding HCl or NaOH to water, what caused the pH
to change or not change as it did? Explain the chemistry that was taking place.
Explain any anomalies.)
buffer resist pH change?)
Results Summary:
1.___ In general, adding a strong acid or a strong base to water _____________
A. neutralizes the solution making a solution that is neither acidic nor basic.
B. drastically changes the pH to be either strongly acidic if strong acid is added, or strongly basic if strong
base is added.
C. causes only a slight change in pH making the solution a only a little more acidic than it already was if
strong acid is added, or only a little more basic than it already was if strong base is added.
D. has variable results as there is no general trend
2.___ In general, adding a strong acid or a strong base to a buffer solution __________
A. neutralizes the solution making a solution that is neither acidic nor basic.
B. drastically changes the pH to be either strongly acidic if strong acid is added, or strongly basic if strong
base is added.
C. causes only a slight change in pH making the solution a only a little more acidic than it already was if
strong acid is added, or only a little more basic than it already was if strong base is added.
D. has variable results as there is no general trend
CH110 Lab 5. Acids and Bases (W14)
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CH110 Lab 5. Acids and Bases (W14)