Use of the pH Meter to Investigate
Acids, Bases and Salts
Learning Goals:
1. Learn to distinguish between strong and weak acids.
2. How to use pH values to determine Equilibrium Constants.
3. Understand the reasons for the pH of a salt solution.
Abstract:
One definition of acids and bases is that they serve as
sources of protons (aka Hydrogen ions), H+, and
hydroxide ions, OH-, respectively. Aqueous solutions of
these compounds can be classified as either strong or
weak, depending on the amount of dissociation. Strong
acids and bases dissociate completely while weak acids
and bases only slightly dissociate. The solutions to be
investigated in this experiment are strong acids and
bases, weak acids and bases, salts of weak acids or
weak bases.
A pH meter can be used to test the pH of prepared
solutions and gain a better understanding of aqueous
solutions of buffers, strong acids and bases, weak acids
and bases and their conjugates.
Your text, Shultz, has an explanation of Equilibrium,
Equilibrium Constants, both Strong and Weak Acids and
Bases and their Salts. See Chapter 8, Pages 270-290.
Prelab Assignment:
In your lab notebook, prepare the following information:
1. A brief (2-3 sentence) introduction to the lab.
2. A table of safety information including the chemicals used in the
lab and their safety handling precautions. This information can
be obtained from the MSDS safety sheets.
The video clips on Using the pH Meter will be helpful. Quick
Time video player may be needed.
Chemicals
Glassware and Supplies
Liquids:
Distilled Water
1.0 M Acetic acid (HC2H3O2)
1.0 M Hydrochloric acid (HCl)
1.0 M Sodium hydroxide (NaOH)
•
•
•
Solids:
•
Sodium acetate,
anhydrous (NaC2H3O2)
•
Ammonium chloride (NH4Cl)
Sodium chloride (NaCl)
•
pH meter with
electrode
standard buffer
solutions pH 4,7,10
100 mL volumetric
flasks
10 mL volumetric
pipets
pipet fillers
magnetic stirrer and
stir bar
Procedures:
A. Distilled Water
1. Boil about 200 mL of distilled and/or deionized
water for five minutes in a beaker and then
cool it to room temperature.
While the distilled water is being heated and cooled you
can set up your pH meter and check that it is working
correctly.
Here is a pH Meter check list:
1. Connect the electrode to the meter.
2. Turn on the pH meter.
3. Place the electrode into the pH 7.0 Buffer
solution.
4. Adjust the meter values to pH 7.0.
5. Check the pH of about 100 ml of room
temperature distilled water.
6. Add 5 drops of 1M HCl to the water stir and
then take a pH reading.
7. Add 10 Drops of 1M NaOH to the water, stir
and then take a pH reading.
Were the pH values of the solutions reasonable?
Were the readings on the meter stable?
Do you think that the values for the solutions were
reproducible? How could you check?
2. Once the water from step (1) above has cooled to
room temperature, add a stir bar and place the
beaker on a magnetic stirrer. Set up the pH meter
as shown below. Make sure the electrode is safely
away from the stir bar and is adequately immersed
in the water.
3. Record the pH and then vigorously swirl the water
with the stirrer for 15 seconds. This will mix air,
which contains carbon dioxide, into the water.
Record the pH again and repeat the stirring and
measurement process at 15-second intervals,
recording each time and pH value, until there is no
further change in the pH.
Carbon dioxide (CO2) readily dissolves in water but
only about 1% of the dissolved CO2 reacts with water
to form Carbonic acid, in the following manner,
CO2(aq) +H2O(l) � H2CO3(aq)
The carbonic acid (H2CO3) is a weak acid and partially
dissociates, Ka1 = 4.2x10-7, causing the pH to
gradually be lowered.
+
H2CO3(aq) � HCO-3(aq) +H(aq)
!"HCO-3 #$ !"H+ #$
K a1 =
[H2CO3 ]
Because of this chemical reaction between water and
the CO2 in the air, pure deionized water does not
necessarily have a pH value equal to pH 7.0. How
would you go about getting the water to return to a
more neutral pH?
B. Strong Acids
Remember that all dilutions must be carried out
using deionized water.
1. Prepare 100 mL of 0.10 M HCl by diluting 1.0 M HCl
using a 10 mL volumetric pipette and 100 mL
volumetric flask. Transfer this solution to a clean dry
labeled beaker.
2. Repeat the above instruction using the previously
prepared solutions to make the following
concentrations of HCl.
a.
b.
c.
d.
0.010 M HCl
0.0010 M HCl
0.00010 M HCl
0.000010 M HCl
3. Record the pH of each solution.
Hydrochloric acid completely dissociates in
water and is by definition a strong acid. Since it
completely dissociates, all of the acid is ionized and is in
the form of H+ ion and the Cl- ion.
+
HCl(aq) ! H(aq)
+ Cl(aq)
Therefore, the hydrogen ion concentration, [H+],
should be equal to the HCl concentration, [HCl], and the pH
should be the negative log of the initial HCl concentration,
which is known.
pH = - log !" H + #$ != - log [ HCl]
C. Weak Acids
1. Using 1.0 M Acetic acid as a starting point, make
five dilutions as in the preceding section to prepare:
a. 0.10 M acetic acid
b. 0.010 M acetic acid
c. 0.0010 M acetic acid
d. 0.00010 M acetic acid
e. 0.000010 M acetic acid
2. Record the pH of each solution.
3. Calculate the Ka for each solution.
+
HOOCCH3(aq) � OOCCH-3(aq)
+ H(aq)
3
!"OOCCH-3 #$ !"H+ #$
Ka =
[HOOCCH3 ]
D. Salts
1. Prepare 100 mL of the following solutions
a. 0.10 M Sodium chloride (NaCl)
b. 0.10 M Ammonium chloride (NH4Cl)
c. 0.10 M Sodium acetate (NaOOCCH3)
2. Determine the pH of each salt solution.
Some salts react with water when they dissolve. This
reaction is called hydrolysis. Cations from strong
bases and anions from strong acids do not undergo
hydrolysis, but cations from weak bases and anions
from weak acids do undergo hydrolysis.
Salts composed of a cation from a strong base (i.e.,
NaOH) and an anion from a strong acid (i.e., HCl),
(e.g., NaCl), dissociate completely when dissolved in
water and do not effect the pH of the solution.
+
NaCl(s) ! Na(aq)
+ Cl(aq)
However, salts composed of a cation from a weak
base (i.e., NH4OH) and an anion from a strong acid
(i.e., HCl), (e.g., NH4Cl), hydrolyze to produce the
weak base and hydrogen ions, which makes the
solution acidic.
NH4Cl(aq) � NH+4(aq) + Cl(aq)
+
NH+4(aq) +!H2O(l) � NH4OH(aq) + H(aq)
Salts composed of an anion from a weak acid (i.e.,
HOOCCH3) and a cation from a strong base (i.e.,
NaOH), (e.g., NaOOCCH3), hydrolyze to produce
the weak acid and hydroxide ions, which makes the
solution basic.
+
NaOOCCH3(aq) ! Na(aq)
+!OOCCH-3(aq)
3
!OOCCH-3(aq)
+!H2O(l) !!HOOCCH3(aq) + OH(aq)
3
Post Lab Assignment
1.
2.
3.
4.
Include the following information in a lab report to give to
your TA.
Graph the pH vs. time data collected in Procedure: A. What is
the final pH obtained by stirring boiled distilled water in the
atmosphere? Determine the mass of CO2 dissolved in the 200
mL of water to obtain this pH.
Using the data collected in Procedure: B make a chart of [HCl],
[H+], measured pH, calculated pH and the numerical difference
between measured and calculated pH. Give reasons for the
differences.
Using the data collected in Procedure: C make a chart of
[HOOCCH3], [H+], measured pH, the Ka determined from the
experimental data, and any numerical difference
between determined Ka and the accepted value found in
the text (Ka=1.8x10-5). Give reasons for any differences.
How are the three salts in Procedure: D different? What is the
origin of the cations and anions in each (strong base, strong
acid, weak base or weak acid)?
Copyright (c) 2012, the ICN Team.