Quantitative Reactions and Titrations Experiment

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CHM 151 Lab 6 – Quantitative Reactions and Titrations

Experiment 7Quantitative Reactions and Titrations

Chemicals: 6 M NaOH, Potassium Hydrogen Phthalate (KHP), sulfuric acid solution ofunknown molarity, phenolphthaleinMaterials: buret, Erlenmeyer flasks, volumetric pipet

Waste Disposal: Completed titration solutions should be disposed of in the red bucketlabeled Acid/Base Waste. Your instructor will neutralize the waste before disposal.

Titrations

In Experiment 4, we learned that the mole is used to determine the quantity of a substance by

one of two methods: gravimetric analysis or volumetric analysis. In this experiment, we willuse volumetric analysis to determine the concentration of an unknown acid solution using abase solution with a known concentration. The reaction between an acid and a base istermed neutralization and involves the combination of hydrogen ions (or hydronium ions, ifyou prefer) with hydroxide ions to form water:

H+(aq) + OH-(aq) H2O(l)

In order to accurately determine the concentration of your solution, you will perform theexperiment in two parts, over two weeks (two lab sessions):

1) Standardization of NaOH: You will prepare a solution of sodium hydroxide with aconcentration of approximately 0.3 M by diluting a stock solution of NaOH. (See theExamples for a sample dilution problem.) You will determine the exact concentrationof the dilute NaOH solution by reacting it with a known amount of the solid organicacid potassium hydrogen phthalate, often abbreviated KHP. The structure, givenbelow, is not terribly important except to note that it is a monoprotic acid meaning thatthere is one mole of H

+ per mole of acid.

H

H

H

H

OH

O

O

O

K

Figure 1: The structure of KHP, KHC8H4O4, Molar Mass = 204.2 g


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The reaction between KHP and NaOH is given below. Notice that there is a one-to-one mole ratio of KHP to NaOH required per the balanced reaction. One mole of KHPwill neutralize one mole of NaOH.

KHC8H4O4(aq) + NaOH(aq) H2O(l) + KNaC8H4O4(aq)

The point in the titration when the number of moles of each reactant are equal iscalled the equivalence point. In other words, # moles NaOH = # moles KHP at theequivalence point.

For a simple one-to-one mole ratio reaction, this point is also called the end point.More precisely, the end point refers to the observed physical change associated withthe equivalence point, usually a color change in a chemical indicator. If done properly,the end point and equivalence point should be very near one another. Therefore, atthe end point (as indicated by a color change in the indicator) you will assume anequivalence point where # moles NaOH = # moles KHP.

Moles of KHP will be calculated from the mass of KHP used and the molar mass forKHP provided above. Thus, if you measure the volume of base, NaOH, required toneutralize a known amount of KHP, you will be able to calculate the molarity of theNaOH solution.

M =moles of NaOH

liters of NaOH

Today, you will use a chemical indicator called phenolphthalein which is colorless ata pH < 8.2 and turns a faint fuscia or pink color at 8.3 and above. You will perform

this reaction in triplicate and calculate an average concentration in molarity of NaOHsolution. The triplicate determinations should agree to within 1%. If they do not,repeat the standardization a fourth time and consult with your instructor.

To calculate the percent agreement:Highest Conc. - Lowest Conc.

Lowest Conc.× 100

2) Concentration of an unknown H2SO4 solution: As a repetition of the same conceptsfrom Part 1, you will calculate the number of moles of H2SO4 in a measured volume ofsolution by titrating to the end point with the standardized NaOH solution.

H2SO4 (aq) + 2 NaOH(aq) 2 H2O(l) + Na2SO4(aq)

Notice that the mole to mole ration between H2SO4 and NaOH is not 1:1. You willneed to perform a stoichiometry calculation to find the moles of H2SO4 from the molesof NaOH required to reach the endpoint of the titration. Once you have found moles ofH2SO4 you can then calculate the molarity of the acid solution by dividing the numberof moles of H2SO4 by the volume of H2SO4 in liters.


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You will perform this titration in triplicate and calculate an average concentration inmolarity of H2SO4 solution. Again, the triplicate determinations should agree to within1%. If they do not, repeat the titration a fourth time and consult with your instructor.

Examples:

Dilution:How will you prepare 500 mL of a 0.2 M solution of NaOH using a 6 M NaOH stock solution?

Solution:When adding solvent to a solution (dilution) the number of moles of solute remain constant.Moles of solute can be found by multiplying the concentration of the solution (Molarity) by thevolume in liters:

moles = M x V (L)

We show that the moles of solute stay the same during a dilution as:

M1V1 = M2V2

You are asked to prepare 500 mL of 0.2 M NaOH so V2 = 500 mL and M2 = 0.2 M. The stocksolution has a concentration of 6 M so M1 = 6 M:

(6 M)V1 = (0.2 M)(500 mL)

V1 = 16.7 mL (volume of 6 M stock solution)

To find the amount of distilled water to add to create 500 mL of the dilute solution:

500 mL total – 16.7 mL stock 6 M NaOH = 483.3 mL of water

Titration Calculations:What is the molarity of the NaOH solution considering 1.050 g of KHP was neutralized by26.45 mL of the NaOH solution?

Solution:Given the reaction: KHC8H4O4(aq) + NaOH(aq) H2O(l) + KNaC8H4O4(aq), you knowthat # moles NaOH = # moles KHP.

moles KHP = 1.050 g KHP 1 mole KHP204.2 g KHP

= 0.0051420 mol KHP = mol NaOH

M =moles of NaOH

liters of NaOH =

0.0051240 mol NaOH

0.02645 liters of NaOH= 0.1937 M NaOH


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Procedure

Preparation:To prepare a buret for use, you should rinse the buret thoroughly with distilled water, thenadd approximately 10-mL of the solution to be used in the buret. (Note: Do Not rinse theburet with the stock NaOH solution! You must prepare your dilute NaOH solution beforepreparing the buret.) Each rinse should flow freely through the stopcock, making sure it doesnot leak at the stopcock and it turns freely. Collect each rinse for disposal. To fill the buret,close the stopcock and add solution to above the top volume mark. Allow the buret to settlefor at least 30 seconds then open the stopcock to drain the solution to below the top volumemark and to release any air bubbles. You should not try to stop the volume exactly on themark. Simply read and document whatever volume is contained.

Reading a buret:Liquids contained in a buret will give a curved meniscus at the surface. As you previouslyhave done, make sure to read the graduations at the bottom of the meniscus. Keep your eye

level with the meniscus when measuring volumes and verify the buret is completely vertical. Itmay be helpful to hold a piece of paper behind the buret when reading the scale. Whendispensing a solution from the buret, remove any hanging drops from the buret tip by gentlytouching it against the side wall of the container. Question for thought: Considering thegraduations on the buret are given to the tenth of a mL (0.1 mL), to how many decimal placesshould you report the volumes on a buret?

Week 1: Standardization of a NaOH solution1. Prepare 500 mL of ~0.3 M NaOH solution using the stock solution provided.

Remember that for a dilution: M1V1 = M2V2.2. Prepare a buret and fill with the NaOH solution to be standardized. Record an initial

volume of NaOH solution.3. Weigh approximately 1.000 g of KHP and transfer into an Erlenmeyer flask, accurately

record the mass of KHP. Be sure to measure at least 1.000 g of KHP so that there are4 significant figures in your mass.

4. Add approximately 50 mL of distilled water to the flask and swirl gently to completelydissolve the KHP. Add two drops of the phenolphthalein indicator to the flask.

5. Slowly add the NaOH solution to a flask while stirring gently. As the NaOH is added,you may see a pink color appear and disappear upon swirly. As the end point isneared the pink color may remain longer. You should then add the NaOH solutiondrop-by-drop until the faint pink color appears and remains after continued swirling forat least 1 minute. This is the end point of the titration. Your goal is to have the faintest

pink color remain at this point. Adding too much NaOH will result in a darker pink colorand will skew your results.

6. Record the final volume of NaOH solution in the buret. Calculate the amount of NaOHrequired to titrate the KHP sample to the end point. Using the volume of NaOHsolution and mass of KHP, calculate the molarity of the NaOH solution.

7. Repeat steps 2-6 for remaining two additional KHP samples. Calculate the averagemolarity of the NaOH solution. If any of the results disagree by greater than 1%,repeat the titration a fourth time and contact your instructor.


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Your report for Week 1 must include the average concentration of your NaOH solution andthe range of the trials used to calculate the average.

Week 2: Titration of H2SO4 with NaOH to determine the Molarity of an H2SO4 solution1. Using a volumetric pipet, dispense exactly 10.00 mL of the H2SO4 solution of unknown

concentration into an Erlenmeyer flask. [Review the procedure for use of a volumetricpipet from Experiment 1 if needed.] Add two drops of the phenolphthalein indicator toeach flask.

2. Following the procedure as in Part 1, slowly titrate the H2SO4 solutions with thestandardized NaOH solution to the end point. Make sure to record the initial volumeand final volume of NaOH for each titration.

3. Calculate the molarity of the H2SO4 solution using the volume of NaOH and its molarityas determined in Part 1. Calculate the average molarity of the H2SO4 solution. If anyof the results disagree by greater than 1%, repeat the titration a fourth time andcontact your instructor.

Your report for Week 2 must include the average concentration of the unknown H2SO4

solution as well as the range of the trials used to calculate the average.

Discussion Questions:

1. Suppose your sample of KHP was inadvertently contaminated with a substance that isneither an acid nor a base. Would this cause your concentration of NaOH to be falselyhigh or falsely low? Justify your answer using the calculations from the experiment.

2. How would the contaminated KHP effect the calculated concentration of the unknownH2SO4 solution, would it be too high or too low? Again, justify your answer using thecalculations from the experiment.

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Quantitative Reactions and Titrations Experiment