CHM 2045L- Equivalent Mass of an Acid 1 Equival.docx

CHM 2045L- Equivalent Mass of an Acid
1
Equivalent Mass of an Acid
Objectives: Upon successful completion of this laboratory the
student will be able to:
1) Perform an acid-base titration accurately to an indicator
endpoint.
2) Calculate moles from molarity and volume.
3) Write the complete, and net ionic equation for the
neutralization of an acid with a base.
4) Calculate equivalent mass from total mass and moles of
hydronium ion.
5) Write a formal scientific communication (laboratory report).
Introduction: Titration is a simple and very frequently used
technique of quantitative volumetric
analysis, which is able to achieve great precision and accuracy
when it is done properly. The titration
apparatus is shown in Figure 1. It consists of a Burette (A), a

clamp (B), a stand (C) and a container (D)
in which the titration reaction occurs. The Burette has a valve
(E) that allows precise control of the flow
of liquid from the burette, and it has a thin tip (F) that produces
small and very uniform drops.
Figure 1, Titration Apparatus
There is a solution that has a very precisely known
concentration of one of the reactants in the Burette.
This is called the titrant. The flask has an unknown amount of
the other reactant, called the analyte. The
analyte can be a known volume of a solution of unknown
concentration, or it can be a carefully weighed
http://chemwiki.ucdavis.edu/Analytical_Chemistry/Quantitative
_Analysis/Titration/Acid-Base_Titrations
https://www.google.com/#safe=off&q=analyte
2
solid compound or mixture dissolved in a solvent. In this lab
you will titrate a solid that you have
weighed to the nearest 1 mg and then dissolved in water.

Notice that the burette is marked the opposite way that a
graduated cylinder is marked. It has the 0 mark
at the top and the 50 mark at the bottom. Rather than being how
much the burette contains, these marks
represent how much has been removed from the burette, if the
level starts at exactly 0.00 ml. What if you
start titrating at some number other than 0? Then simply
subtract your initial measurement from the final
measurement.
When you are reading a burette, just as with any other
instrument, your measurement precision should go
1 decimal place past the smallest tic mark (Figure 2)
Figure 2 How to read a burette
The smallest tic mark on our burettes is 0.1 ml. This means that
you will read the burettes to the nearest
0.01 ml.
3

The number of moles of analyte present can be determined
easily from the volume of the titrant, the
concentration of the titrant in
��
��
(molarity, also abbreviated as M), and the balanced equation of
the
reaction between the titrant and the analyte. This can tell you a
number of things. If you know the
formula and molar mass of the analyte, it can tell you how many
grams are present, and the percent
composition of a carefully weighed sample of the analyte
material. If the volume of the analyte solution
is known precisely, it can tell you the molar concentration of
the analyte solution, and if you know the
mass of a pure unknown compound, it can tell you the molar
mass, of that compound.
The equation to obtain the moles of analyte from the volume of
titrant is as follows:
� �� (��) ×
1�
1000��
×
��

1 � ��
×
��
��
= ��
Equation 1 calculation of moles of analyte
To find the mass, multiply by molar mass
�� ×
��
1 ��
= ��
Equation 2, calculation of mass of analyte
To find percent composition, divide by total grams mixture:
��
��
× 100% = ��
Equation 3, Calculation of percent composition of an analyte
mixture
To find the molar mass of an analyte divide the mass of the
analyte by the moles of analyte
��
��
=

��
1 �� ⁄
Equation 4 calculation af the molar mass of an analyte
Understanding the experiment: In this experiment we will
perform an acid – base neutralization reaction,
and use equations 1 and 4 to find the equivalent mass of an
acid. Acids are broadly defined as sources of
hydrogen ions, H+, also called protons. In water, acids will
react with water to form hydronium ions,
H3O+, by the following reaction, where HA stand for a generic
acid. This reaction is also called acid
dissociation.
�� + �2� → �3�
+ + �−
Equation 5 dissociation of a monoprotic acid in water
In fact, the proton, or H+ ion, never exists alone in a water
solution. It always exists as the hydronium
ion, H3O+. Often people will talk about the hydrogen ion and
refer to it as H+, but what they really mean is
hydronium ion.
Acids can be mono-protic, di-protic, tri-protic and even poly-
protic, depending on how many hydrogen

ions they can donate. The stoichiometric ratio of a mono-protic
acid is 1 to 1; that of a diprotic acid is 1
to 2; that of a triprotic acid is 1 to 3 and so forth. The reaction
of a strong di-protic acid with water would
be:
4
�2� + 2 �2� → 2�3�
+ + �2−
Equation 6, Dissociation of a diprotic acid in water
Notice the stoichiometric ratio of hydronium ion to acid is 2 to
1. The equivalent mass of the acid is the
amount required to produce 1 mole of hydronium ion. It would
take half as many moles of the acid in
equation 2 to make a mole of hydronium ion as it would if it
were a monoprotic acid.
From this, you can see that the equivalent mass of a monoprotic
acid will be equal to its molar mass,

while the equivalent mass of a diprotic acid will be ½ of its
molar mass. For a triprotic acid it would be
1/3 and so forth. Examples of monoprotic, diprotic and triprotic
organic acids are shown in Figure 3. The
acidic proton is shown in bold.
Figure 3 Organic acid structures.
A base is broadly defined as a compound that absorbs hydrogen
ions. Bases produce hydroxide ions, OH-
, in water in one of two ways. They either dissociate in water to
form hydroxide ions (These are called
Arrhenius bases), or they react with water to produce hydroxide
ions. The base that we will use in this
laboratory, sodium hydroxide, is one of the ones that dissociates
in water. The equation is below.
��(�)
�2�
→ ��− (��) + ��+(��)
Equation 7, Dissociation of an Arrhenius base in water
Observe that sodium hydroxide will produce exactly as many
moles of hydroxide ion as there are moles
of sodium hydroxide that dissolve. Ammonia is an example of a
base that react with water to form
hydroxide ion. These bases are called Brønsted-Lowry bases.

The equation for the reaction of ammonia
is shown below:
��3 + �2� → ��
− + ��4
+
Equation 8, Dissociation of a Brønsted Lowry base in water
A major simplification that is being made in this description of
acids and bases is the assumption that they
dissociate or react completely with the water to form hydroxide
or hydronium ions. While this is true of
strong acids and bases, there are many weak acids and bases
that only react a little bit before the reaction
starts going in the other direction to establish what is called an
equilibrium with only a very low
concentration of hydronium or hydroxide ion. A complete
description of weak acids and weak bases is
beyond the scope of this course. You will study this and other
aspects of equilibrium in grueling detail, in
CHM 2046. No worries, though; you will have a whole lot more
chemistry under your belt by then.
Even though all of the acids that will be used in this laboratory
are considered weak acids, they will
completely dissociate through the course of the titration,
because the sodium hydroxide is a strong base,

and it will completely react with the small amount of hydronium
produced by any aqueous acid, no matter
http://chem-guide.blogspot.com/2010/03/concept-of-equivalent-
mass.html
http://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Ba
ses/Acid/Arrhenius_Concept_of_Acids_and_Bases
ses/Acid/Bronsted_Concept_of_Acids_and_Bases
http://www.chem1.com/acad/webtext/chemeq/
5
how weak it is. This will drive more of the acid to dissociate
and make more hydronium ion, which will
in turn be gobbled up by the hydroxide ion, until there is no
acid left. This tendency is known as Le
Chatelier's principle. It is also a topic that will be covered
extensively in CHM 2046.
The equations of acid and base add together as follows
�� (��) + �2�(�) → �3�
+(��) + �−(��)
+
��(�)

�2�
→ ��− (��) + ��+(��)
+
�3�
+ + ��− → 2�2�(�)
=
��(�) + �� (��)
�2�
→ �2�(�) + ��
+(��) + �−(��)
scheme 1, Reaction of an acid and a base
The third equation is called the net ionic equation for acid base
neutralization. It can be derived by
assuming that the acid and the base are present in their
completely dissociated forms.
�3�
+(��) + �−(��) + ��− (��) + ��+(��)
�2�
→ �2�(�) + ��
+(��) + �−(��)
Equation 9, Total ionic equation of an acid base reaction
The ions that are crossed out are called spectator ions, because
they appear on both sides of the arrow.
Taking them out gives you the third equation in scheme 1.

You can Also see that the stoichiometric ratio for a dibasic acid
is two to one, base to acid, and for a
tribasic acid the stoichiometric ratio of base to acid is 3 to 1, as
shown in equations 10 and 11.
2��(�) + �2� (��)
�2�
→ 2�2�(�) + 2��
+(��) + �2−(��)
Equation 10 Reaction of a diprotic acid with sodium hydroxide
3��(�) + �3� (��)
�2�
→ 3�2�(�) + 2��
+(��) + �3−(��)
Equation 11, reaction of a triprotic acid with sodium hydroxide
Because 1 mole sodium hydroxide reacts with 1 mole hydronium
ion, the equivalent mass of the acid is
the mass of the acid divided by the moles of sodium hydroxide.
In other words:
�� (
�
��⁄ ) =
�� (�)
��
×
1000��

1 �
×
1
��
�⁄ �� ℎ��
In most cases, both reactants and products of acid base reactions
are colorless. It would therefore be
impossible to see when the reaction is complete. To determine
this we need to add an indicator dye.
Indicator dyes are dyes that react with something in the reaction
mixture to change color when the
reaction is done. We will use dye molecule called
phenolphthalein, which is a very weak acid that is
much less likely to give up its protons than the acids that we are
titrating. When phenolphthalein does
give up its protons, it turns pink, or red. When the very last
molecule of the acid reacts, there is no more
hydronium ion to react. This is called the equivalence point.
When the equivalence point is reached, the
hydroxide ion in the next drop of titrant will react with the
phenolphthalein and turn it red.
https://www.khanacademy.org/science/chemistry/chemical-
equilibrium/factors-that-affect-chemical-equilibrium/v/le-

chatelier-s-principle
https://www.khanacademy.org/science/chemistry/chemical-
equilibrium/factors-that-affect-chemical-equilibrium/v/le-
chatelier-s-principle
ses/Case_Studies/Acid_and_Base_Indicators
https://en.wikipedia.org/wiki/Phenolphthalein
https://en.wikipedia.org/wiki/Equivalence_point
6
Figure 4 structure of phenolphthalein acidic hydrogens are
shown in bold
This marks the endpoint of the titration. At the true end point,
very little phenolphthalein will have
reacted, so your solution will be a very light pink. If it turns
dark pink, you will have added too much
base. See Figure 5. The flask on the left is a perfect endpoint.
The one on the right has too much base
added.
Figure 5, Good endpoint (left) overshot endpoint(right)
It is important to continuously swirl your analyte solution. If
you do not, you can get a false endpoint.

The color will appear, but then disappear when you stir it. As
you approach the endpoint clouds of pink
color will appear briefly when you add the base, then disappear
(figure 6).
Figure 6, transient pink cloud near endpoint
7
Procedure:
1) Obtain a vial of unknown acid from the chemical stockroom.
2) Obtain the following equipment: burette clamp, ring stand,
Burette with valve and tip, burette funnel, 3
clean 250ml Erlenmeyer flasks, a 250 ml beaker, 2 or 3 little
squares of white paper, a squirt bottle, and
a few plastic transfer pipettes. Make sure that the valve fits
snugly in the burette and that the tip fits
snugly in the valve.
3) Wash out the squirt bottle with deionized water and fill it

with deionized water. Then wash the burette,
two of the three flasks and the beaker with deionized water.
Dry the beaker with a clean paper towel.
4) Dispense about 150 ml of the sodium hydroxide solution
from the carboy into the beaker. Write down
the molar concentration of this solution.
Caution! Sodium hydroxide is very caustic and it will
permanently blind you if it gets in your
eyes, even in low concentrations. Wear approved Safety glasses
or goggles!
5) Assemble the burette in the burette clamp, and use a transfer
pipette to run a few pipettes full of the
sodium hydroxide solution down the inside walls of the burette.
Put the unwashed Erlenmeyer flask
under the burette, and drain out the sodium hydroxide solution
into the Erlenmeyer flask. Repeat this
process 2 more times.
6) Place the funnel in the top of the burette and carefully pour
the sodium hydroxide until it reaches close to
the 0.00 ml mark.
7) Open the valve and let a few drops of the sodium hydroxide
titrant run into the waste flask. This will fill
the tip of the burette with titrant.

8) Discard the waste solution in the sink and wash the flask
thoroughly with deionized water.
9) Take the unknown sample of acid to the balance. Put a plastic
weigh boat onto the balance and press
“tare”. When the balance reads 0.000g, weigh out the amount of
unknown acid that is indicated on the
vial to the nearest 0.001g. Do not exceed this amount, or you
might not be able to titrate it with only 1
burette full of sodium hydroxide solution. Write the mass down
on your data sheet.
10) Carefully pour the acid powder into one of the flasks. Use
the corner of the weigh boat to pour from.
With your squirt bottle, wash any solid that remains on the
weigh boat into the flask. Mark this flask
“rough”
11) Put about 50 ml of deionized water into the flask and then
swirl the flask to dissolve as much of the acid
as possible. Add a few drops of phenolphthalein solution to the
flask.
12) Use a ruler or the edge of a notebook as a straight edge, and
draw a thick dark line horizontally across
one of the small pieces of white paper. Place the other piece
under the burette, and place the flask with

the acid solution on top of it.
13) Hold the paper with the line behind the burette, so that the
line is horizontal, just underneath the
meniscus. This will reflect off of the meniscus, making it
easier to read (see Figure 7). With your eye at
the level of the meniscus, read the burette to the nearest 0.01
ml.
14) While constantly swirling the flask of analyte, open the
valve and rapidly titrate until the acid solution
turns pink. Be ready to stop the flow when the color change
occurs.
15) Write the initial volume, final volume and net volume (final
– initial) in the “rough” section of the data
sheet.
16) Refill the burette with sodium hydroxide solution and let a
little run through the pipette tip into the
titrated sample if it is necessary to refill the tip of the burette.
17) Repeat steps 9 through 11 with the other two flasks. Mark
them “trial 1” and “trial 2”.

8
18) To help you to estimate the amount of titrant that will be
needed for trials 1 and 2, you can do a
proportional calculation as shown below. This will allow you to
titrate quickly to just under the
estimated volume, and then titrate slowly to get an accurate
endpoint.
�� =
��ℎ ��
��ℎ ��
× ��
19) Read the initial volume as in step 13, and add the estimated
net volume for Trial 1 to the initial volume
to get the estimated final volume.
20) Titrate rapidly to about 5 ml before the estimated final
volume for Trial 1. Then titrate the solution drop
by drop, with constant swirling until it turns a very light pink.
21) Measure the final volume as in step 13, and write down the
measured initial volume and final volume in
the “Trial 1” column of your data sheet.
22) Refill the burette and repeat steps 19 through 21 for trial 2.

23) Calculate the equivalent mass of the acid for trials 1 and 2,
then calculate the average value.
24) Discard the titrated acid solutions and the excess sodium
hydroxide solution in the sink with water.
Clean and return all the equipment. Return the acid sample to
the stockroom, and clean up your work
area.
Figure 7, reading the meniscus with a black line
9
Report Sheet: Equivalent Mass of an Acid
Name:
_____________________________________________________
______________________
Lab
Partner(s):____________________________________________
_________________________
Class period: ______________________________________

Date: __________________
Data sheet: to be turned in only with full, formal lab report.
Unknown number
NaOH molarity
Rough: used to estimate endpoint for titrations in trials 1 and 2
Acid Mass Initial Volume Final Volume Net Volume
Trial 1 Trial 2
Acid mass (g)
Estimate the net volume
needed for this mass of acid
based on the rough
Your Initial volume (ml)
What is your estimated final
volume (ml)

Your Measured final
volume (ml)
Your Measured net volume
(ml)
Measured net volume (L)
Moles OH─
Moles H3O+
Equivalent mass
of acid (g/mol)

Average equivalent mass ____________________
10
Prelab: Equivalent Mass of an Acid
Name:
_____________________________________________________
______________________
Class period: ______________________________________
Date: __________________
Show calculations and be mindful of significant figures for full
credit.
The following data were observed in an equivalent mass of an
acid experiment.
1) Fill in the blanks (2 points each). Show all calculations for
full credit.
NaOH molarity Acid mass
Initial volume
(ml)
Final volume
(ml)

Net volume
(ml)
0.1000 M 0.2511g
Numerical value:
2) (3 points) How many moles of hydroxide ion were consumed
in the titration?_________________
3) (3 points) How many moles of hydronium ion were available
from the acid?___________________
4) (4 points) What is the equivalent molar mass of the acid?
____________
5) (4 points) If it happened that the acid in this experiment was
one of the ones represented
in the table below, what is the most likely identity of this acid?
___________________
Acid name Molar mass Number of protons

Butanoic acid 88.11 g/mol 1
Tartaric acid 150.087 g/mol 2
Citric acid 192.124 g/mol 3
11
Formal Laboratory Report
This lab requires a formal laboratory report that will be turned
in online through Turnitin. Specific
guidelines for writing the report are shown below:
Section Requirements
Introduction
(10 points)
be how the
stoichiometry of the
acid base reaction can be used to volumetrically determine the
equivalent mass of

the acid.
industry, or
environmental protection.
for example: “The volume and
concentration of the base
solution are used to find the number of moles of acid present.”
not “I will use the
volume and concentration of the base solution to find the
number of moles of acid
present.”
ences with sufficient detail that your instructor
can find them.
Procedure
(20 points)
procedure in the lab
manual
of equal experience
can duplicate the experiment
filled to 1 cm above
the 0 ml mark with a 0.097 M sodium hydroxide solution.” not
“Fill the burette to

1 cm above the 0 ml mark with a 0.097 M sodium hydroxide
solution.”
Data and
calculation
(60 points)
graphs, and also described
in paragraph form.
sses
should be recorded
to within 0.001g.
acid, and equivalent
mass of the acid.
Results and
discussion
(30 points)
equivalent mass of the acid
for all titrations, and
the average values.

evaluate how closely they
agree.
highest equivalent masses
in this experiment are about 200 g/mol, and the lowest
equivalent mass organic
acid is oxalic acid, with an equivalent mass of 45 g/mol.
Anything much less than
this is probably not reasonable, and masses of more than 500
are also not
reasonable in this experiment.
Conclusion
(10 points)
reliability of your results
results of this
experiment, discuss whether or not you could perform the
titration in the practical
example that you provided in the introduction with sufficient
precision and
accuracy
Prelab

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