CHM 2045L- Equivalent Mass of an Acid

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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

CHM 2045L- Equivalent Mass of an Acid

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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.

CHM 2045L- Equivalent Mass of an Acid

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The number of moles of analyte present can be determined easily from the volume of the titrant, the

concentration of the titrant in 𝑚𝑜𝑙𝑒𝑠

𝑙𝑖𝑡𝑒𝑟 (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:

𝑉 𝑡𝑖𝑡𝑟𝑎𝑛𝑡 (𝑚𝑙) × 1𝐿

1000𝑚𝑙 × 𝑚𝑜𝑙𝑒𝑠 𝑡𝑖𝑡𝑟𝑎𝑛𝑡

1 𝐿 𝑡𝑖𝑡𝑟𝑎𝑛𝑡 × 𝑚𝑜𝑙𝑒𝑠 𝑎𝑛𝑎𝑙𝑦𝑡𝑒

𝑚𝑜𝑙𝑒𝑠 𝑡𝑖𝑡𝑟𝑎𝑛𝑡 = 𝑚𝑜𝑙𝑒𝑠 𝑎𝑛𝑎𝑙𝑦𝑡𝑒

Equation 1 calculation of moles of analyte

To find the mass, multiply by molar mass

𝑚𝑜𝑙𝑒𝑠 𝑎𝑛𝑎𝑙𝑦𝑡𝑒 × 𝑔𝑟𝑎𝑚𝑠 𝑎𝑛𝑎𝑙𝑦𝑡𝑒

1 𝑚𝑜𝑙𝑒 𝑎𝑛𝑎𝑙𝑦𝑡𝑒 = 𝑔𝑟𝑎𝑚𝑠 𝑎𝑛𝑎𝑙𝑦𝑡𝑒

Equation 2, calculation of mass of analyte

To find percent composition, divide by total grams mixture:

𝑔𝑟𝑎𝑚𝑠 𝑎𝑛𝑎𝑙𝑦𝑡𝑒

𝑔𝑟𝑎𝑚𝑠 𝑚𝑖𝑥𝑡𝑢𝑟𝑒 × 100% = 𝑝𝑒𝑟𝑐𝑒𝑛𝑡 𝑎𝑛𝑎𝑙𝑦𝑡𝑒

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

𝑔𝑟𝑎𝑚𝑠 𝑎𝑛𝑎𝑙𝑦𝑡𝑒

𝑚𝑜𝑙𝑒𝑠 𝑎𝑛𝑎𝑙𝑦𝑡𝑒 = 𝑔𝑟𝑎𝑚𝑠

1 𝑚𝑜𝑙𝑒 𝑎𝑛𝑎𝑙𝑦𝑡𝑒⁄

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.

𝐻𝐴 + 𝐻2𝑂 → 𝐻3𝑂 + + 𝐴−

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:

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𝐻2𝐴 + 2 𝐻2𝑂 → 2𝐻3𝑂 + + 𝐴2−

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.

𝑁𝑎𝑂𝐻(𝑠) 𝐻2𝑂 → 𝑂𝐻− (𝑎𝑞) + 𝑁𝑎+(𝑎𝑞)

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:

𝑁𝐻3 + 𝐻2𝑂 → 𝑂𝐻 − + 𝑁𝐻4

+ 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

CHM 2045L- Equivalent Mass of an Acid

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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

𝐻𝐴 (𝑎𝑞) + 𝐻2𝑂(𝑙) → 𝐻3𝑂 +(𝑎𝑞) + 𝐴−(𝑎𝑞)

+

𝑁𝑎𝑂𝐻(𝑠) 𝐻2𝑂 → 𝑂𝐻− (𝑎𝑞) + 𝑁𝑎+(𝑎𝑞)

+

𝐻3𝑂 + + 𝑂𝐻− → 2𝐻2𝑂(𝑙)

=

𝑁𝑎𝑂𝐻(𝑠) + 𝐻𝐴 (𝑎𝑞) 𝐻2𝑂 → 𝐻2𝑂(𝑙) + 𝑁𝑎

+(𝑎𝑞) + 𝐴−(𝑎𝑞) 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.

𝐻3𝑂 +(𝑎𝑞) + 𝐴−(𝑎𝑞) + 𝑂𝐻− (𝑎𝑞) + 𝑁𝑎+(𝑎𝑞)

𝐻2𝑂 → 𝐻2𝑂(𝑙) + 𝑁𝑎

+(𝑎𝑞) + 𝐴−(𝑎𝑞) 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.

2𝑁𝑎𝑂𝐻(𝑠) + 𝐻2𝐴 (𝑎𝑞) 𝐻2𝑂 → 2𝐻2𝑂(𝑙) + 2𝑁𝑎

+(𝑎𝑞) + 𝐴2−(𝑎𝑞) Equation 10 Reaction of a diprotic acid with sodium hydroxide

3𝑁𝑎𝑂𝐻(𝑠) + 𝐻3𝐴 (𝑎𝑞) 𝐻2𝑂 → 3𝐻2𝑂(𝑙) + 2𝑁𝑎

+(𝑎𝑞) + 𝐴3−(𝑎𝑞) 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:

𝐸𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡 𝑚𝑎𝑠𝑠( 𝑔 𝑚𝑜𝑙⁄ ) =

𝑀𝑎𝑠𝑠 𝑎𝑐𝑖𝑑 (𝑔)

𝑚𝑙 𝑁𝑎𝑂𝐻 × 1000𝑚𝑙

1 𝐿 ×

1

𝑚𝑜𝑙𝑒𝑠 𝐿⁄ 𝑠𝑜𝑑𝑖𝑢𝑚 ℎ𝑦𝑑𝑟𝑜𝑥𝑖𝑑𝑒

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.

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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

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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”.

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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.

𝑒𝑠𝑡𝑖𝑚𝑎𝑡𝑒𝑑 𝑡𝑟𝑖𝑎𝑙 𝑛𝑒𝑡 𝑣𝑜𝑙𝑢𝑚𝑒 = 𝑟𝑜𝑢𝑔ℎ 𝑛𝑒𝑡 𝑣𝑜𝑙𝑢𝑚𝑒

𝑟𝑜𝑢𝑔ℎ 𝑎𝑐𝑖𝑑 𝑚𝑎𝑠𝑠 × 𝑡𝑟𝑖𝑎𝑙 𝑎𝑐𝑖𝑑 𝑚𝑎𝑠𝑠

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

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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 ____________________

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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

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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)  Explain the objective of the experiment and describe how the stoichiometry of the

acid base reaction can be used to volumetrically determine the equivalent mass of

the acid.

 Provide an example of how titration is used in medicine, industry, or environmental protection.

 Write in passive voice 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.”

 Cite any references with sufficient detail that your instructor can find them.

Procedure

(20 points)  Write the procedure in your own words. Do not copy the procedure in the lab

manual

 The procedure should contain sufficient detail that a chemist of equal experience can duplicate the experiment

 Use passive voice past tense. For example: “The burette was 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.”

 Cite any references

Data and

calculation

(60 points)

 All quantitative results should be presented as tables or graphs, and also described in paragraph form.

 Volumes should be recorded to the 0.01 ml place, and masses should be recorded to within 0.001g.

 Use Passive voice past tense

 Show the calculations for net volume, moles base, moles acid, and equivalent mass of the acid.

Results and

discussion

(30 points)

 List the values obtained for the equivalent mass of the acid for all titrations, and the average values.

 Compare the values of each titration to each other, and evaluate how closely they agree.

 Discuss whether or not your results are reasonable. The 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.

 Use passive voice in this section as well.

Conclusion

(10 points)  Discuss any experimental factors that could influence the reliability of your results

 Given the equipment provided and your evaluation of the 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

(20 points)