CHEMISTRY 161 – EXPERIMENT 9
ACIDS, BASES AND SALTS: pH Measurement and Neutralization
INTRODUCTION
Acids are compounds that donate protons, H+ (actually H3O+), in solution. Acids have a sour taste and are
involved in various chemical reactions. They react with many metals to form a salt and hydrogen gas; they react
with carbonates and bicarbonates to form a salt, water and carbon dioxide; they react with bases to form a salt and
water. This last type of reaction is called neutralization.
Bases are compounds that react with acids by accepting protons. Bases include not only hydroxides, OH-,
but also oxides, carbonates, bicarbonates and ammonia. The word alkali is often used to mean a base.
Salts are compounds which produce positive and negative ions when dissolved in water. They are called
electrolytes and are often formed by the reaction of an acid with a base.
Acids and bases may be classified according to their degree of ionization in solution. Strong acids and bases
are highly ionized, whereas weak acids and bases are only slightly ionized. Common strong acids include nitric,
sulfuric,, hydrochloric, perchloric, hydrobromic and hydroiodic. Most other acids are weak. Most hydroxide bases
are strong, except for ammonium hydroxide, which is weak.
The degree of ionization of an acid or base may be determined by measuring the pH of the solution.
Mathematically pH is defined as:
pH = – log [H+]
Where [H+] is the concentration of hydrogen ions, protons, in the units moles per liter. The table below lists pH
values and their corresponding [H+] and [OH-].
pH
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
H+
100
10-1
10-2
10-3
10-4
10-5
10-6
10-7
10-8
10-9
10-10
10-11
10-12
10-13
10-14
OH-
10-14
10-13
10-12
10-11
10-10
10-9
10-8
10-7
10-6
10-5
10-4
10-3
10-2
10-1
100
At pH 7, the [H+] and the [OH-] are equal; the solution is neutral. Any solution with a pH less than 7 is
acidic; the lower the pH, the more acidic the solution. Solutions with a pH greater than 7 are basic; the higher the
pH, the more basic the solution.
Large pH differences exist in the body. The gastric juice is very acidic having a pH of about 1. The pH of
the intestinal contents varies from 2 to 6. Saliva is nearly neutral, its pH being near 7. Urine is slightly acidic, with a
pH range of 4.5 to 7. Blood is slightly basic, with a pH of about 7.4. If the pH of the blood falls below 7.1 or rises
above 7.8, death ensues.
pH is very important in many areas of the medical and biological sciences. Bacteria grow only in culture
media of certain pH; changes in pH may inhibit their growth or even destroy them. Tissue cultures must be
maintained at pH values corresponding to those of body fluids or they will not survive. Most drugs are weak acids
or bases and are present in the body both the ionized and nonionized forms. The proportion of ionized to
nonionized drug depends upon the pH of the solution in which the drug is dissolved. The nonionized molecules
are more fat soluble and can diffuse more easily across cell membranes; ionized molecules are not fat soluble and
are less able to diffuse across most body membranes.
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Aspirin, for example, is a weak acid. It is absorbed into the blood stream, principally from the stomach and
upper portion of the small intestine. In both of these areas the pH is low. At a low pH, aspirin does not ionize
and, therefore, the amount of nonionized aspirin is high. To treat an aspirin overdose, the physician must convert
the nonionized aspirin to its ionized form in order to slow down its absorption into the blood stream. Since aspirin
is a weak acid, an alkali such as sodium bicarbonate is used to convert aspirin to its sodium salt, which is ionized.
Salt hydrolysis is the reaction between water and the ions in a salt that produces a change in the pH. If a salt
is derived from a weak acid and a strong base (sodium bicarbonate), hydrolysis will produce a basic solution. If a
salt is derived from a weak base and a strong acid (ammonium chloride), hydrolysis will produce an acidic solution.
If a salt is derived from an acid and a base of equal strength, the solution will be neutral (sodium chloride). Salts
frequently are used to change the pH values of a variety of solutions: medications, culture media, cleaning agents
and swimming pools. Salts also play an important part in buffer solutions which you will examine in the next
experiment.
An easily demonstrated acid-base reaction that takes place in the stomach is the neutralization of excess
stomach acid by antacid tablets. Excess stomach acid can cause indigestion and heartburn. When an advertisement
claims that an antacid tablet consumes 47 times its own weight in excess stomach acid, it means that 1 gram of that
product will neutralize 47 grams of 0.1 M hydrochloric acid. You will check this claim by performing an acid-base
titration as described in this experiment.
PROCEDURE
PART I: pH Measurements
Reminder: Do the first step of Part II and let the solutions sit before you begin Part I.
1. Place about 3 mL of each of the various solutions listed below into clean test tubes. Obtain a pH paper kit
containing 1 roll of wide range pH paper (Alkacid Test Ribbon) and 6 rolls of short range alkacid paper.
2. You will determine the pH of each solution in a two step process. First, rinse a stirring rod thoroughly
with deionized water, dip it into one of the solutions and touch it to a piece of wide range pH paper.
Compare the color of the pH paper to the color chart on the side of the roll and record the wide range pH
on the Report Sheet.
3. Second, select one of the short range pH papers base on the information from the wide range
measurement. Use the stirring rod again to transfer a drop of solution to the short range paper and
compare the color to the chart on the side of that short range roll. Record your measurement.
4. Rinse the stirring rod thoroughly with deionized water and repeat the entire procedure for each of the
remaining solutions.
PART II: Neutralization Titration
1. Accurately weigh ½ of a Rolaids tablet to the nearest 0.01 g and record the information on the Report
Sheet.
2. Place the tablet in a 250 mL Erlenmeyer flask and add exactly 100 mL of 0.1 M HCl.
3. Repeat this procedure with a second tablet, being sure to label each flask.
4. Occasionally swirl each flask to help the tablets to dissolve while you do part A of this experiment.
5. Obtain a buret and fill it with 0.1 M NaOH, making sure that all bubbles have been removed from the tip
of the buret.
6. After the tablet has dissolved as completely as possible, add 5-7 drops of bromphenol blue indicator to the
flask.
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7. Record the initial volume of 0.1 M NaOH in the buret and titrate until the solution turns from yellow to
just barely blue.
8. Record the final volume of NaOH in the buret.
9. Refill the buret and repeat this titration for the second tablet.
CALCULATIONS:
The way the experiment is designed, an excess of 0.1 M HCl is added to the Rolaids tablet. The tablet
is completely neutralized and the remaining HCL is determined by the titration. Since the concentration of
the NaOH solution is equal to the concentration of the HCl solution, the volume of HCL remaining is
numerically equal to the volume of NaOH used. Assume 1 mL of 0.1 M HCl weighs 1 gram, then the grams
of HCl consumed equals the mL of HCl consumed.
POST-LAB QUESTIONS:
1. Why is the pH of deionized H2O often less than 7.0?
2. Decide whether the salts listed below should yield an acidic, basic, neutral or indeterminate pH when
dissolved in deionized H2O. Give chemical reactions to support your answer.
NaCl, NaHCO3, NaC2H3O2, NH4C2H3O2 and NH4Cl
3. How do the theoretical pH ranges determined in question 2 compare with your experimental results?
Explain any discrepancies.
4. Calculate your average value for acid consumed by a Rolaids tablet in your two trials and compare your
average to the manufacturer's claim?
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Name_______________________________
REPORT SHEET
ACIDS, BASES AND SALTS: pH Measurement and Neutralization
PART I.
pH Measurements
Solution
Wide Range pH
Short Range pH
Deionized Water
_____________
_____________
0.1 M HCl – Hydrochloric Acid
_____________
_____________
0.1 M NaOH – Sodium Hydroxide
_____________
_____________
0.1 M NaCl – Sodium Chloride
_____________
_____________
0.1 M NaHCO3 – Sodium Bicarbonate
_____________
_____________
0.1 M NaC2H3O2 – Sodium Acetate
_____________
_____________
0.1 M NH4C2H3O2 – Ammonium Acetate
_____________
_____________
0.1 M NH4Cl – Ammonium Chloride
_____________
_____________
5% Acetic Acid – Vinegar
_____________
_____________
1% Ivory Liquid
_____________
_____________
1% Calgon
_____________
_____________
Soda
_____________
_____________
Milk
_____________
_____________
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PART II.
Titration
Trial 1
Trial 2
Weight of Rolaid
___________
___________
Final volume of NaOH
___________
___________
Initial volume of NaOH
___________
___________
Volume of NaOH used
___________
___________
100.0 mL
100.0 mL
Volume of HCl remaining
___________
___________
Volume of HCl consumed
___________
___________
Grams of HCl consumed
___________
___________
Grams of HCl consumed/gram Rolaid
___________
___________
Initial volume of HCl
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