CHMI 1006E - Lab5 Titration - Fall2020 PDF

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Lab 5 Titrations

LAB #5: TITRATION: PERCENT MASS OF CALCIUM CARBONATE

KEY CONCEPTS

        

Titration curves Primary Standard Neutralization reaction Strong/weak base/acid Indicators Titrant Analyte Precision Accuracy

INTRODUCTION

A titration is a procedure used to determine a concentration of an acid or a base. During this procedure,

a solution

of

known concentration (a

titrant)

is

added

to a

solution of

unknown

concentration. In this experiment, oxalic acid H2C2O4 will be a primary standard to calibrate a titrant of NaOH.

0

Base (ex. NaOH) 10

20

30

40

50

Acid

Note: NaOH is a strong base and dissociates as follows:

NaOH

(s)

→

+

Na

(aq)

+

OH

(aq)

Once a base is added to an acid, the following neutralization reaction occur:

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Lab 5 Titrations

For H2C2O4:

2OH

(aq)

+ H2C2O4

(aq)

→

2-

C2O4

(aq)

+

2 H 2O

(l)

For The medicinal ingredient in Tums tablets is CaCO3. HCl will be used to neutralize the calcium carbonate as follow:

CaCO3 (s) + H

-

HCO3

(aq)

+ H

+ (aq)

+

↔

(aq)

2+

↔

Ca

H2CO3 (aq)

-

(aq)

↔

+ HCO3

(aq)

CO2 (g) + H2O

(l)

By adding an excess amount of acid (HCl), all the carbonate will be neutralized and an excess of HCl will remain in solution. This excess of HCl will be titrated with NaOH. Therefore the amount of acid neutralized by the Tums tablet can be calculated. This approach is called back-titration.

Both neutralization reactions produce water. Since the concentration of H

𝑝𝐻 = −𝑙𝑜𝑔[𝐻3 𝑂

solution decreases, the pH of the solution increases (

+]

+

ions present in the

).

In this experiment, you will add 0.1 M NaOH to an acid until the solution’s colour changes from clear to pink. The solution will change its colour because it contains phenolphthalein (which is an acidbase indicator). This indicator is colourless when pH is 0-8.2 and pink when pH is 8.2-12. It is imperative to stop adding base at the initial onset of pink apparition to ensure your endpoint is as close to the actual equivalence point as possible.

Once the solution is pink in colour, the reaction has reached the equivalence point. All of the acid -

present in the flask has reacted with NaOH. The number of moles of OH added to the flask is equal to the number of moles of H

+

that was originally present in the flask.

𝑛𝐻 + = 𝑛𝑂𝐻 −

At equivalence point:

𝑛𝐻 + = number Since 𝑛 = 𝐶 × 𝑉 Where

of

moles

of

H

+

(mol)

and

𝑛𝑂𝐻 − =

number

of

moles

of

OH

-

(mol).

𝑛𝐻 + = 𝐶𝑂𝐻 − 𝑉𝑂𝐻− where,

𝐶𝑂𝐻 − =

-

concentration of OH (M), and

𝑉𝑂𝐻 − =

volume of OH

-

(L). The volume of OH

-

-

is the

-

amount of OH added to the solution using the burette and the concentration of the OH is known from the standardization. The moles of H from the moles of H

+

+

are easily solved. The moles of the acid can be determined

using the following ratios:

𝑛𝐻 + = 𝑛𝑎𝑐𝑖𝑑

,

when the acid is monoprotic

𝑛𝐻 + = 2𝑛𝑎𝑐𝑖𝑑

,

𝑛𝐻 + = 3𝑛𝑎𝑐𝑖𝑑

,

when the acid is diprotic

when the acid is triprotic

(example: HCl)

(example: H2SO4)

(example: H3PO4)

The molar concentration of the acid can be calculated by dividing the moles of the acid by the volume of acid used in the flask.

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Lab 5 Titrations

A titration curve illustrates the change in pH when a titrant is added to a flask.

14

12

10

The equivalence point +

-

([H ]=[OH ])

8

Hp 6

4

2

0 0

5

10

15

20

25

30

35

40

Volume of titrant added (mL)

A titration begins with a slow rise in pH. OH

-

+

ions react with H , but the solution is being diluted

(the volume is increasing). At the equivalence point, all of the acid has reacted with the OH

-

ions.

The pH rises quickly since there is no more acid in the solution to lower the pH. A titration ends with -

a slow rise in pH. More OH ions are added to the solution, but the solution is being diluted (there is no more acid to react with and the volume is increasing).

OBJECTIVES

1.

To prepare a NaOH solution and a standard solution of H2C2O4.

2.

To standardize the base solution.

3.

To titrate a Tums tablet and calculate the percent mass of calcium carbonate.

From the titration you will find the number of moles of acid in excess that did not react with calcium carbonate present in the solution. Knowing this the number of moles of calcium carbonate can be calculated. You will then find the mass of calcium carbonate in the tablet using n=m/molar mass (where n= number of moles (mol), m= mass (g), molar mass (g/mol)). You will then use the following equation to calculate the mass %.

Mass % =

CHMI 1006

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mass of calcium carbonate total mass of the tablet

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Lab 5 Titrations

PROCEDURE

PART A: SOLUTION PREPARATION

1.

Calculate the mass, in grams, of solid NaOH pellets required to have a 500 mL solution of NaOH, 0.080 M.

2.

Show the calculation on the report form.

Weigh the NaOH and record the exact mass of NaOH on the report form. Prepare the NaOH in the polyethylene bottle.

Add deionized water to about half the bottle.

NOTE The exact volume of water is not important. The concentration should be approximately 0.080 M. Totally dissolve the NaOH by inverting the capped bottle several times, then proceed to the next step. Keep the bottle sealed at all times when not in use.

3.

∙

Calculate the mass of H2C2O4 2H2O, in grams, required to make a 100 mL solution of H2C2O4, 0.12 M.

4.

Show the calculation on the report form.

∙

Prepare the 0.12 M solution of H2C2O4 2H2O in the 100 mL volumetric flask.

∙

Record the

∙

exact mass of H2C2O4 2H2O and calculate the actual concentration of the H 2C2O4 2H2O solution on the report form.

NOTE

Refer to step 2 of Lab #2 for detailed instructions on how to properly use the volumetric flask to prepare accurate solutions.

The volumetric flask has 5 significant figures.

PART B: STANDARDIZATION OF THE NaOH

1.

Prepare the burette and fill it with the NaOH (titrant) solution.

NOTE

Do not attempt to set the initial volume exactly to zero! Screw the lid tightly on the bottle of NaOH when it is not in use.

2.

3.

Using the proper technique, pipette 10.00 mL of 0.12 M H2C2O4 into an Erlenmeyer flask.

Add approximately

50 mL of deionized H2O to

the Erlenmeyer and three drops of

phenolphthalein indicator solution.

4.

Before beginning the titration, estimate the volume of NaOH required to titrate the analyte. Show the calculation on the report form.

5.

Record the initial volume of NaOH from your buret te into Table 5.1.

Dispense the NaOH

until you have added 5 mL less than your estimation.

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NOTE: From this point forward, add the NaOH in a drop wise fashion while constantly swirling your flask. Place a the white side of the contrast card under the flask, to help distinguish the appearance of the deprotonated indicator solution (pink color).

6.

Once the analyte solution is light pink, record the final burette reading in Table 5.1 of the report form.

7.

Repeat the titration 2 more times.

8.

Determine if your titrations are precise enough to continue. If the variability in the volume of titrant used, for your three trials, is:

a.

Within ± 0.1 mL with respect to the volume of titrant used (Not from your estimate), move on to part C.

b.

Is less then ± 0.25 mL, complete another trial in order to obtain three trials within the 0.1 mL limit.

c.

Greater than ± 0.25 mL, consult with a TA or the laboratory technologist before continuing.

PART C: TITRATION OF TUMS TABLETS

1.

Obtain three Tums tablets from your laboratory technologist. Weigh each tablet and record their masses on the report form.

1.1. Collect 80 mL of 0.5 M HCl into a 100 mL graduated cylinder

1.2. Pour the HCl into a 150 mL beaker

2.

Using the proper technique, pipet 25.00 mL of standardized 0.5 M HCl into an Erlenmeyer. Record the actual concentration of the HCl used in the report form.

3.

Add one antacid tablet and note any chemical reactions.

4.

After the tablet is completely dissolved, add about 25 mL of water and heat until it is close to boiling.

5.

Add three drops of phenolphthalein indicator solution to the sample flask.

6.

Record the initial burette reading in Table 5.2 of the report form.

7.

Perform a titration on the sample to the equivalence point. Record the final burette reading in Table 5.2 of the report form.

8.

Repeat steps 2 to 7 with the two other samples.

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LAB #5: TITRATIONS - REPORT FORM To be filled by the student

Student Name:

Station Number:

Student Number:

Date:

Section Day:

Weekday

day

/month /year

Section:

#

To be filled by a TA or the laboratory instructor

⎕ ⎕

Safety Violation: (If applicable, -10%)

PPE Chemical

⎕ ⎕

Behavior Equipment

Late Submission: (If applicable)

dd

/

mm

/

yy

Time

Comment

⎕ ⎕

Initials

-25% -100%

Initials

Report Sign Off:

Initials

PART A:

1) Calculation of the mass of NaOH required for the solution.

Calculated mass of NaOH:

2)

Actual mass of NaOH:

3)

Calculation of the mass of H2C2O4 2H2O required.

______________

________________

∙

∙

Calculated mass of H2C2O4 2H2O:

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____________

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Lab 5 Titrations

4)

∙

Actual mass of H2C2O4 2H2O:

___________

Calculate the actual concentration of H2C2O4 from the actual mass.

Actual concentration of H2C2O4 in solution:

___________

PART B:

4)

Estimate volume of base required to complete the titration.

-

The equivalence point for this titration will occur when molOH = molH the moles of reactants (H

Note:

+

+

. Note that you are comparing

-

and OH ).

1 mol of H2 C2O4 will react with 2 moles of OH

-

DO NOT use C1 V1 = C2 V2

VNaOH:

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Lab 5 Titrations

Table 5.1: Volume of NaOH used to titrate 10.00 mL of 0.12 M H2C2O4

Trial #1

Trial #2

Trial #3

Trial #4

Trial #5

Trial #6

Final burette reading

Initial burette reading

Total volume delivered

Average of best three trials

PART C:

1) Mass of tablets

Sample 1: ___________

Sample 2: ___________

Sample 3: ___________

2) Actual concentration of HCl supplied: __________________

Table 5.2: Volume of NaOH used.

Trial #1

Trial #2

Trial #3

Final burette reading

Initial burette reading

Total volume added

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QUESTIONS

1)

Calculate the molar concentration, with correct number of significant figures, of the NaOH solution from the data tabulated in Table 5.1.

Clearly show your calculations below.

Molar concentration of NaOH:

2)

___________

Using the data in Table 5.2 and the mass of the tablets, calculate the mass and percent mass of calcium carbonate in each tablet. Clearly show your calculations for one trial below. Perform the calculation for each sample individually. Include your answers in Table 5.3.

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(Calculation continued)

Table 5.3: Final results from Tums titration

Mass of calcium carbonate per tablet

Percent calcium carbonate per tablet

Sample 1

Sample 2

Sample 3

3a.) Explain your observations when you dropped a tablet into HCl. What is the gas produced? (Show a chemical reaction)

3b.) How many moles of that gas should be produced? (Assume 500 mg of calcium carbonate)

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GENERAL QUESTIONS

Adipic acid (H2C6H8O4) is a carboxylic acid containing compound of industrial importance, notably due to its use during fabrication of nylon.

Adipic acid can be titrated using NaOH by the following

reaction:

H2C6H8O4 + 2 NaOH

→2H O + Na C H O 2

2

6

8

4

a) Suppose you titrate 0.283 g of pure adipic acid, what volume (in mL) of 0.1000 M NaOH solution would be required to reach the equivalence point?

b) Suppose you do the same titration with 0.304 g of industrial grade, 74% (w/w) adipic acid.

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(General questions continued)

Name the following compounds:

-

OH :

∙

H2C2O4 2H2O: 2-

C2O4 :

NaCl:

NaOH:

H2O: +

H :

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