Acids, Bases, and pH
Michael Jones, Ph.D.
Collin College
Department of Chemistry
There are several definitions of acids and bases, but the most commonly-used definition in chemistry is
the Bronsted-Lowry definition. In this system, an acid is a proton (H +) donor, and a base is a proton
(H+) acceptor. For example:
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
In this reaction, HCl donated a proton (H+) to NaOH, or more specifically, to the OH -, thus forming
water as a product. In addition to OH-, other common bases in water often contain CO32-, or HCO3-.
Both of these bases accept H+ from acids to form carbonic acid, which tends to immediately
decompose when formed under such conditions.
H2CO3(aq)  H2O(l) + CO2(g)
We often categorize acids as “strong” or “weak” based upon how much they dissociate (or how easily
they give up their protons). HCl(aq) is a strong acid because it dissociates completely in water:
HCl(aq)  H+(aq) + Cl-(aq)
But other common acids such as acetic acid or phosphoric acid only dissociate to a small extent (that
is, they are not willing to give up their protons; they prefer their “molecular” form, on the left). Notice
the “equilibrium arrows” used to denote the lack of complete dissociation.
CH3COO-(aq) + H+(aq)
CH3COOH(aq)
H3PO4(aq)
H+(aq) + H2PO42-(aq)
For a more quantitative picture of acid strength, we often turn to the pH scale. This measure of acid
strength is defined as the negative log (base 10 log) of the H+ concentration.
pH = - log [H+]
Two acids of the same concentration (specifically molar concentration or molarity) may have different
pH values, depending on if they are strong or weak, and dissociate to a different extent. Common pH
values range from 0 to 14, with low values (0-2) representing strongly acidic solutions, high values
(12-14) representing strongly basic solutions, and a pH=7 representing neutral solutions, that is, neither
acidic nor basic.
We will also look briefly at buffer solutions. Buffer solutions are a mixture of a weak acid and its salt
or a weak base and its salt. They are called buffer solutions because they “buffer” or level out large pH
changes upon addition of acid or base to a solution. Blood contains several different buffers in order to
maintain the pH within a narrow range.
PROCEDURE:
I.
Add ~ 2 mL of each of the following solutions (individually) to 5 small test tubes, and test
the pH of each of the following solutions: 1.0 M HCl(aq), vinegar, ammonia, 1.0 M
NaOH(aq), and tap water, using both pH paper and the hand held pH meters.
To use the pH paper, tear off about a ¼ to ½ inch strip from the roll, dip a glass stirring rod
into the test solution, and touch the glass rod to the strip of pH paper. Compare the color to
the color chart on the roll and assign a pH value to each solution. Rinse the glass rod with
deionized water and wipe dry between each solution.
Remove the protective cap from the hand-held pH meters (unless this has already been
done) and place in the same solutions above. These pH meters are generally calibrated
before use to respond to different pH values. Again, rinse the pH meter with deionized
water before and in between testing each of the 5 solutions. Record the values in your
table, give the pH meter a final rinse and return it.
II.
Antacids are common OTC medication used to neutralize “excess acid”. We would not
want to completely neutralize all stomach acid, since a low pH is necessary for the enzymes
that aid in digestion to function. Common antacids may contain CaCO3, NaHCO3, and/or
Mg(OH)2. We will use 1.0 M HCl to simulate stomach acid.
Weigh out 0.100 g of each of the 3 solids above and place in 3 separate test tubes. Add 1.0
M HCl to each of the test tubes, 1 drop at a time (count the number of drops), until each
solid dissolves completely and/or stops “fizzing”.
III.
Place 4 mL of deionized water in a clean test tube, and 4 mL of pH 4 buffer in another
clean test tube. Check the pH of the 2 samples, using pH paper as described in procedure I.
Add 1 drop of 1.0 M HCl(aq) to each test tube and check the pH again. Add a 2 nd drop of
1.0 M HCl(aq) and check the pH one more time. Make sure you stir the samples well after
adding each drop, before testing the pH.
ACIDS, BASES, and pH POSTLAB
I.
pH of various solutions:
a. Solution
pH with pH paper
pH with pH meter
1.0 M HCl
_______
________
Vinegar
_______
________
Ammonia
_______
________
1.0 M NaOH
_______
________
Tap water
_______
________
b. Do your pH values determined by 2 different methods generally agree with each
other?
c. We did not test the following household products, but predict (or find out)
whether the following would be acidic or basic.
Drain cleaner
___________________
Lemon juice _________________
II.
Soft drink ____________________
Bleach ______________________
Antacids
Mass
drops of 1.0 M HCl
CaCO3
__________
______________
NaHCO3
__________
______________
Mg(OH)2
__________
______________
Should it take the same number of drops to neutralize the same mass of these
substances? Explain.
Extra strength antacids may contain 750 mg of CaCO3 in a single tablet. If we
assume 20 drops per mL, how many mL of 1.0 M HCl would one of these tablets
neutralize (based on your data above)?
III.
Buffers
Initial pH
pH after 1 drop HCl
pH after 2 drops HCl
Deionized water
______
________
________
pH 4 buffer
______
________
________
What was the initial pH difference between the 2 samples?
What was the pH difference after 1 drop of HCl? 2 drops?
Did the buffer do what a buffer should do?