EXPERIMENT 19: Properties and Reactions of Acids-Bases
Chemicals:
Materials:
6M NaOH
6M HC2H3O2
Phenolphthalein indicator
Universal indicator
0.1M HCl
0.1MNaOH
0.1M HC2H3O2
0.1M NH3
1M HCl
1M NaOH
1M HC2H3O2
6M HCl
Plastic pipette dropper
Test tubes
Stirring rod
pH paper
Tooth pick
Watch glass
Hot plate
Thermometer
INTRODUCTION- Traditionally, solutions were labeled as being acidic or basic based on their taste and
texture. Those that tasted sour were said to be acidic and solutions that tasted bitter and were slippery to touch
were said to be basic. Thus, substances such as lemon juice and vinegar were identified as acids, and solutions
of lye and caustic soda as bases. Several definitions have been proposed for acids and bases. Depending upon
the situation, one or more definition is applicable. The three acid-base theories are the Arrhenius theory, the
Bronsted-Lowry theory, and the Lewis theory.
The Arrhenius theory applies to solutes in aqueous solution. Arrhenius acids are solutes that ionize in aqueous
solution to produce hydrogen ions, H+ (hydronium ion, H3O+). Examples include hydrochloric acid, sulfuric
acid, etc. Hydrochloric acid ionizes in aqueous solution in one step as follows :
HCl(aq)  H+(aq) + Cl(aq)
or
HCl(aq) + H2O(l)  H3O+(aq) + Cl(aq)
Sulfuric acid, a diprotic acid, ionizes in two steps;
H2SO4 (aq)  H+ (aq) + HSO4 (aq)
HSO4 (aq)  H+ (aq) + SO42 (aq)
Those solutes in aqueous solution that ionize or dissociate to produce hydroxide ions, OH are called Arrhenius
bases.
KOH (aq)  K+ (aq) + OH (aq)
Johannes Bronsted and Thomas Lowry expanded the definition of an acid to include all compounds that can
be proton donors and the definition of a base to include all compounds that are proton acceptors. This definition
is not restricted only to solutes based on their behavior in aqueous solution and hence is applicable to wider
class of compounds than the Arrhenius definitions. In this theory, the ionization of an acid in water is viewed as
a proton transfer reaction where the proton acceptor is a water molecule and hence a Bronsted-Lowry base. The
products that result from an acid-base reaction in this theory are called the conjugate base of the acid and the
conjugate acid of the base.
1
For example, consider again the ionization of hydrochloric acid in aqueous solution :
HCl (aq) + H2O (l)  H3O+ (aq) + Cl (aq)
Acid
base
conjugate
conjugate
acid
base
The conjugate acid is the species formed when a Bronsted-Lowry base accepts a proton from a Bronsted-Lowry
acid. The conjugate base is the species that remains after the Bronsted-Lowry acid has lost a proton.
The Lewis definition of acids and bases due to G. N. Lewis is the broadest of the three theories. The Lewis
theory defines an acid as an electron-pair acceptor and a base as an electron-pair donor. This definition includes
reactions that contain neither hydrogen nor hydroxide ions.
BF3 + ‫ ׃‬NH3  F3B‫׃‬NH3
acid
base
(e-pair acceptor)
(e-pair donor)
In the context of the Arrhenius theory, water is termed as being neutral. The Bronsted-Lowry definition
classifies water as an amphoteric species, one capable of behaving both as an acid and as a base. To understand
this, consider the following dissociation of water called autoinization of water;
H2O (l) + H2O (l)
H3O+(aq) + H+ (aq)
In this case, the hydronium ion, H3O+, forms when a proton, H+, is transferred from one H2O molecule to
another. The other species that result from this process is the hydroxide ion, OH-. Thus while one water
molecule, the proton acceptor, functions as the base, the other plays the role of the acid. The ability of one water
molecule to accept a proton from another water molecule or any acid is due to the two lone pairs of electrons on
the oxygen atom of water. The autoionization results in equal molar amounts of H3O+ ions and OH ions and
hence the solution is neutral. In a sample of pure H2O, the concentrations of H3O+ and OH ions at 25 oC are
1.0 x 107 M.
[H3O+] = [OH‾] = 1.0 x 107 M
The concentration of hydronium ion , [H3O+], of a solution is commonly expressed in terms of the pH of the
solution, which is defined as the negative logarithm of [H3O+] or negative logarithm of [H+] in the solution;
pH = – log [H3O+] or pH = – log [H+]
2
Thus the hydrogen ion concentration can be obtained from the pH of the solution as follows;
[H3O+] = 10 – pH
Similarly the concentration of hydroxide ion , [OH-], of a solution is commonly expressed in terms of the pOH
of the solution, which is defined as the negative logarithm of [OH-]
pOH = – log [OH ‾]
The hydroxide ion concentration can be obtained from the pOH of the solution using the equation
[OH‾] = 10 ‾ pOH
Additonally, the pH and the pOH of any aqueous solution are related as are the hydrogen and the hydroxide ion
concentrations. The relevant equations are;
[H+] [OH‾] = 1.0 x 10 –14
pH + pOH = 14
Example (1)
What is the pH and pOH of a solution that contains 3.50 x 10-5 M hydronium ions?
pH = – log [H3O+] = – log (3.50 x 105) = 4.46
pOH = 14 – pH = 14 – 4.46 = 9.54
Example (2)
Calculate the hydronium ion and hydroxide ion concentrations of a solution that has a pOH of 4.40.
pH = 14 – pOH = 14 – 4.40 = 9.60
[H3O +] = 10 – pH = 10 –
9.60
= 2.51 x 1010 M
[OH] = 10 – pOH = 10 – 4.40 = 3.98 x 10 –5 M
In water, [H3O+] is equal to 1.0 x 107 M, so the pH is 7.0. Because [H3O+] = [OH] in water, which is neither
acid nor base.
pH = 7.0 ( neutral)
pH < 7.0 (acidic) pH > 7.0 (basic)
The pH scale has a range of 0.0 to 14.0. A practical way to evaluate the relative acidity or basicity of solutions
is to compare their effect on indicators.
3
Indicators, in chemistry, are natural or synthetic substances that change color in response to the nature of
chemical environment. Litmus, for example, is a natural dye that turns red in most acidic solutions and blue in
most basic solutions. Compounds that undergo color changes when there is a pH change in the solutions in
which they are contained are called indicators (Table 1).
Indicators are used to provide information about the degree of acidity of a substance or the state of some
chemical reactions within a solution being tested or analyzed.
Table 1 – Indicators and pH ranges
Indicators
Acid(color change)
Base(color change)
pH range
Methyl orange
Red
Yellow
3.1-4.4
Methyl Red
Red
Yellow
4.2-6.3
Bromothymol Blue
Yellow
Blue
6.0-7.8
Phenol Red
Yellow
Re
6.4-8.0
Phenolphthalein
Colorless
Pink
8.0-9.8
Thymol Blue
Red
Yellow
1.2-2.8
Alizarin Yellow
Yellow
Red
10.1-12.0
In this experiment, you will observe the properties of acids and bases with suitable indicators. By the end of this
experiment, you will be able to determine the pH of various solutions by observing the color of several
indicators in these solutions and also using pH paper and pH meter.
4
PROCEDURE:
I. Light-bulb Conductivity Test of Acid-Base Solutions: (Instructor Demo)
Your instructor will demonstrate the relative degrees of ionization between the strong and weak acids and
bases by measuring their conductivity. The stronger the acid or base (electrolyte), the brighter the light bulb
glows when the electrodes are placed in the solution. Conversely, the light bulb will glow dimly when
immersed in solutions of weak acid or base.
II. pH of Acid-Base Solutions:
 Obtain 10 drops (~ 0.5 mL) of 0.1 M HCl (aq) (hydrochloric acid) and
0.1 M HC2H3O2 (acetic acid) in two small test tubes.
 Dip your stirring rod into each acid solution and touch the tip on a small piece of pH paper.
Contrasts the color produced to the colors on the
pH-box, then read and record the corresponding pH value.
 Obtain 10 drops (~ 0.5 mL) of 0.1 M NaOH(aq) and 0.1 M NH3 (aq) (aqueous ammonia) in two small
test tubes.
 Dip your stirring rod into each base solution and touch the tip on a small piece of pH paper.
Contrasts the color produced to the colors on the
pH-box, then read and record the corresponding pH value.
III. Production of Salts from Acid-Base Reactions:
a. Combine 20 drops (~ 1 mL) of 6 M HCl and 6 M NaOH on a clean, dry watch glass. Stir with a
toothpick.
b. Place the watch glass on a hot plate at a low setting (3 or 4), and let the solution evaporate to
dryness. Write your observations.
c.
IV. Neutralization
1) Neutralization of 1.0 M HCl with 1.0 NaOH :
a) Phenolthalein Indicator
 Add 20 drops (~ 1 mL) of 1 M HCl (aq) in a large test tube.
 Add 2 drops of phenolthalein indicator. Observe and record the color of the solution.
 Dropwise add 1 M NaOH (aq) till the color of the acid solution changes, while shaking the test tube
in between the drops.
 Count the # of drops of base used for complete neutralization.
 Record the # of drops and the color observed.
 Determine the mL of base used for neutralization and record
(20 drops ≈ 1 mL)
b) Universal Indicator
 Add 20 drops (~ 1 mL) of 1 M HCl (aq) in a large test tube.
 Add 2 drops of universal indicator. Observe and record the color of the solution.
5
 Dropwise add 1 M NaOH (aq) till the color of the acid solution changes, while shaking the test tube
in between the drops.
 Count the # of drops of base used for complete neutralization.
 Record the # of drops and the color observed.
 Determine the mL of base used for neutralization and record
2) Neutralization of 1.0 M HC2H3O2 with 1.0 NaOH :
a) Phenolthalein Indicator
 Add 20 drops (~ 1 mL) of 1 M HC2H3O2 (aq) in a large test tube.
 Add 2 drops of phenolthalein indicator. Observe and record the color of the solution.
 Dropwise add 1 M NaOH (aq) till the color of the acid solution changes, while shaking the test tube
in between the drops.
 Count the # of drops of base used for complete neutralization.
 Record the # of drops and the color observed.
 Determine the mL of base used for neutralization and record
b) Universal Indicator
 Add 20 drops (~ 1 mL) of 1 M HC2H3O2 (aq) in a large test tube.
 Add 2 drops of universal indicator. Observe and record the color of the solution.
 Dropwise add 1 M NaOH (aq) till the color of the acid solution changes, while shaking the test tube
in between the drops.
 Count the # of drops of base used for complete neutralization.
 Record the # of drops and the color observed.
 Determine the mL of base used for neutralization and record
V. Heat of Neutralization:
(CAUTION: Be very careful with these concentrated solutions of acids and bases in this part of the
experiment. They are very damaging to eyes, skin and clothing. Safety glasses are a MUST and labapron is recommended.)
 Obtain four small test tubes and place them in your test tube rack.
 Transfer 2 mL of 6 M HCl into one of the large test tube. Determine and record the solution’s temperature
by use of a digital thermometer.
 Transfer 2 mL of 6 M NaOH into the other large test tube.
 While the test tube with 6 M HCl is still on the rack, swiftly add the
6 M NaOH solution while stirring the mixture with your digital thermometer.
 Observe the temperature change, and record the highest temperature before it starts to decline.
 Repeat the same process this time using 2 mL of 6 M HC2H3O2 (Acetic Acid) and the same base solution, 2
mL of 6 M NaOH. Record your measurements and your observations.
6
EXPERIMENT 19: Properties and Reactions of Acids-Bases
REPORT FORM
Name ___________________________
Instructor ________________________
Date ____________________________
Partner’s Name: ___________________
Results and Observations
I. & II. Conductivity and pH of Acid/Base Solutions:
Solution
Strong or Weak
Conductivity
Strong or Weak
pH
Acid/Base
0.1 M HCl (aq)
0.1 M HC2H3O2 (aq)
0.1 M NaOH (aq)
0.1 M NH3 (aq)
Question 1. Both acids above have the same concentration, 0.1 M, but they exhibit different pH value.
Explain.
Question 2. Both bases above have the same concentration, 0.1 M, but pH of NH3(aq) is lower than that of
NaOH. Explain.
7
III.
a) Observation of HCl/NaOH product.
b) Identity of the product:
c) Write the reaction occurred between HCl and NaOH.
__________________________________________________________________________
IV.
Monitoring Neutralization Reactions by Use of Different Indicators:
20 Drops of
1.0 M Acid
Color Of Acid Solution
After Adding
Phenolthalein
# Of Drops Of Base
Added For Complete
Neutralization
mL of Base
Added
The Color of Solution
At the Point of Complete
Neutralization
Color of Acid Solution
After Adding Universal
Indicator
# Of Drops Of Base
Added For Complete
Neutralization
mL of Base
Added
The Color of Solution
At the Point of Complete
Neutralization
HCl
HC2H3O2
20 Drops of
1.0 M Acid
HCl
HC2H3O2
V. Heat of Neutralization:
Acid
Initial T, oC
Highest T , oC after
adding NaOH
HCl
HC2H3O2
8
Is reaction exo- or
endothermic?
EXPERIMENT 19: Properties and Reactions of Acids-Bases
Name ___________________________
Pre- laboratory Questions and Exercises
Due before lab begins. Answer in space provided.
1. Define neutralization.
2. Give two examples of balanced chemical equations for neutralization reaction. (other than
HCl and NaOH)
3. Provide definitions for the following terms;
a) Bronsted-Lowry acid and base
b) Acidic, basic, and neutral solutions
4. Write three properties of acids and bases.
5. Calculate the [H3O+], hydronium ion concentration, of the followings;
a) pOH = 12.8
d) [OH –] = 3.98 x10 – 4 M
b) pH = 6.65
9
EXPERIMENT 19: Properties and Reactions of Acids-Bases
Name: ___________________________
Post- laboratory Questions and Exercises
Due after completing lab. Answer in space provided.
1. Write the balanced chemical equation between HC2H3O2 and NaOH.
2. Supposedly sulfuric acid, H2SO4, is used to react with NaOH in part III,
what would be the residue once the water is evaporated?
1. Provide definitions for the followings;
pH 
pH paper –
pH meter 
indicator 
3. In this experiment, how do you determine if a salt behaving as an acid or a base?
4. The pH of a solution is measured to be 6.80. Calculate the following;
a) [H3O+]
d) [OH]
b) pOH
10