Experiment #10

Volumetric Analysis

The Titration of Acids and Bases

Purpose

This experiment will allow you to gain practical experience in the preparing standard solutions, using a pipette and a buret, and performing standard titrations. You will use this experience to experimentally determine the concentration of acetic acid in an unknown concentration solution. There are three parts to this experiment, they are

1. Preparation a standard solution of oxalic acid (~0.07 M)
2. Preparation a sodium hydroxide solution (~0.1M) which will be standardized using the standard oxalic acid solution.
3. Determination of the concentration of acetic acid in an unknown vinegar solution. (0.1-0.2M)

Introduction

Standard solutions are solutions with known concentrations, generally to four significant figures. There are two different ways to make a standard solution. We can make a primary or a secondary standard. A primary standard is prepared directly by dissolving a known mass of sample to make a known volume of solution. A secondary standard is prepared by dissolving an approximate amount of sample into a volume of solvent and determining its exact concentration through titration experiments. Primary standards are prepared from compounds that are at least 99.9% pure, have a definite composition, are water soluble, are easily weighed, and do not change composition on contact with air. Oxalic acid dihydrate (H₂C₂O₄•2H₂O) fits these criteria and therefore may be used as a primary standard. Sodium hydroxide absorbs water when it comes into contact with air and therefore it is difficult to obtain a pure, dry sample to weigh. For this reason the sodium hydroxide solution will be titrated with the oxalic acid standard to become a secondary standard.

In the first part of this experiment you will prepare a solution of known concentration (to 4 significant figures) of oxalic acid. The oxalic acid crystallizes with two water molecules per oxalic acid in the crystalline network. For this reason, we will weigh out an appropriate amount of oxalic acid dihydrate to dissolve in water. The water molecules in the crystal network will become part of the water of solution once it is dissolved. For this reason the molar concentration of oxalic acid dihydrate will be the same as the molar concentration of oxalic acid.

In order to standardize the sodium hydroxide solution you will perform a titration. Sodium hydroxide reacts with oxalic acid according to the reaction below:

H₂C₂O₄(aq) + 2 NaOH(aq) à 2 H₂O(l) + Na₂C₂O₄(aq)

You will measure a 25.00 mL aliquot of the oxalic acid solution into a flask and add an indicator. An indicator is a substance that changes color when a solution changes from acidic to basic. The common indicator used for acid base titrations is phenolphthalein. Phenolphthalein is colorless in a solution that is acidic and bright pink in a solution that is basic. In this titration the oxalic acid solution is acidic and therefore phenolphthalein will be colorless. The sodium hydroxide solution will be added drop wise from a buret into the flask containing the oxalic acid and indicator. As the sodium hydroxide is added to the flask it will react with the oxalic acid and be neutralized. At the point where all of the oxalic acid is reacted, the next drop of sodium hydroxide will make the entire solution basic and it will turn pink. At this point you have completed the titration.

In order to determine the concentration of acetic acid in the vinegar solution you will titrate it with the standardized sodium hydroxide solution. The equation for this reaction is

HC₂H₃O₂(aq) + NaOH(aq) à H₂O(l) + NaC₂H₃O₂(aq)

In order to get the best precision possible, you should repeat each titration until you get 3 trial that are within 1% of each other.

Procedure

Safety Notes

Ø Wear safety glasses at all times.

Ø All solutions can go down the drain.

Preparation of Standard Oxalic Acid Solution

Carefully weigh a 100 or 150 mL beaker and record its mass. Measure between 2.1 and 2.3 g of pure oxalic acid dihydrate crystals into the beaker and weigh again. Add 30-60 mL of deionized water to the beaker and dissolve the crystals. You may gently heat the solution to speed up this process. Transfer the solution quantitatively into a clean 250-mL volumetric flask. Rinse the beaker with 15-20 mL of deionized water and pour this solution into the volumetric flask and repeat. This will ensure that all of the oxalic acid is transferred into the volumetric flask. Fill the volumetric flask to within about 2 cm of the mark and allow it to sit for a minute. This will allow any water clinging to the edges of the neck to drain into the flask. Using an eyedropper, fill the flask to the mark with water. Stopper the flask and mix the solution by repeated inversion and swirling. This requires about 30 inversions and takes close to 1 minute. Next you will transfer this solution into one of the clean 500 mL bottles in your drawer. To transfer, first pour a small amount (~10 mL) of the standard solution into the bottle and rinse the inside walls of the bottle. Discard this wash solution. Do this two times to prevent the standard solution from being diluted with any pure water that might have remained in the bottle after washing.

Calculate the molarity of the oxalic acid solution using the mass of acid used and the volume of the volumetric flask and record the concentration of your solution on the bottle. Your label should look like the one pictured in figure 10.2.

Preparation of Sodium Hydroxide Solution

Calculate the volume of concentrated sodium hydroxide solution you must use to prepare 500 mL of approximately 0.1-0.15 M sodium hydroxide solution. (Check the label on the reagent bottle to determine its approximate concentration. The concentration will be 6M) Measure an appropriate amount using your graduated cylinder. Try to measure to within 1 mL of the desired amount of reagent. The exact amount is not important because you will be standardizing this solution later. Pour the concentrated base solution into a clean (need not be dry) 500 mL bottle and fill the bottle up to the shoulder with deionized water. Shake the bottle well and label it as above recording the concentration to 1 significant figure.

Titration

1. Pipette 25.00 mL of oxalic acid solution into a clean but not necessarily dry Erlenmeyer flask. (You may do three samples as you will be doing at least 3 titrations.)

Add 2-3 drops of phenolphthalein to the Erlenmeyer flask. (Do not forget this step or you will not see any endpoint.)

2. Set up a buret using your sodium hydroxide solution as the titrant. Titrate your oxalic acid solution to a pink phenolphthalein endpoint as described in the box. Repeat until you have three titrations that differ from the average (median?) by no more than 0.05 mL of titrant.

3. Pipette 25.00 mL of vinegar solution into a clean but not necessarily dry Erlenmeyer flask and titrate this sample. Add 2-3 drops of phenolphthalein and repeat until you get three titrations that differ from the average by no more than 0.05 mL of titrant.

4. Calculate the molarity of your standard sodium hydroxide solution and your unknown acetic acid (vinegar) solution and complete the excel spreadsheet for this experiment (the spreadsheet can be found on the chemistry department home page, under lab archives). You should be able to determine the concentrations of both the sodium hydroxide and vinegar solutions to 4 significant figures

5. Calculate the molarity of your standard sodium hydroxide solution and your unknown acetic acid (vinegar) solution and complete the excel spreadsheet for this experiment (the spreadsheet can be found on the chemistry department home page, under lab archives). You should be able to determine the concentrations of both the sodium hydroxide and vinegar solutions to 4 significant figures

Prelab Exercise

1. You have weighed out precisely 2.471 g of oxalic acid dehydrate, and diluted it to 250.0 mL. What is its molarity?

2. What are volumetric pipets and burets and what are they used for?

3. To what volume are you to record the values from the buret?

4. What is the name of the indicator you are to use and what color does it turn at the end point?

5. You have diluted 9 mL of nominally 6M NaOH to 500 mL. What is the approximate molarity of the base? (Pay attention to significant figures)

6. A 20.00 mL sample of the standard oxalic acid solution (from 1 above) requires 24.16 mL of your base to reach the end point. Write the equation for the reaction.

7. Calculate the molarity of the base in question #6.

8. 24.37 mL of an unknown acid is titrated with 40.62 mL of the standard base (from #7 above). What is the molarity of this acid?

9. What is the meaning of the term “end point” in a titration and what is true about the system at the end point?

Experiment 10 – Volumetric Analysis

The Titration of Acids and Bases

Data Sheet

Record all of your original data on this sheet in ink and you may transfer it to the lab book pages in pencil to do the work-up.

Preparation of Oxalic acid Standard

Mass of oxalic acid and beaker

Mass of Beaker

Mass of oxalic acid

Concentration of Oxalic acid solution

Sample Calculations

Preparation of Sodium Hydroxide Standard

Molarity of concentrated NaOH stock solution

Volume of NaOH stock solution used

Final volume of dilute NaOH solution

Approximate concentration of NaOH solution

Sample Calculations

Standardization of NaOH Solution

Pipet volume _______________

	Trial 1	Trial 2	Trial 3	Trial 4
Volume of oxalic acid solution used (mL)
NaOH buret, final volume reading (mL)
NaOH buret, initial reading (mL)
Volume of NaOH solution used (mL)
Molarity of NaOH (M)

Average Molarity of NaOH solution______________________

Sample Calculations

Standardization of Unknown Acid

	Trial 1	Trial 2	Trial 3	Trial 4
Volume of unknown acid solution used (mL)
NaOH buret, final volume reading (mL)
NaOH buret, initial reading (mL)
Volume of NaOH solution used (mL)
Molarity of unknown acid (M)

Average Molarity of vinegar solution______________________

Sample Calculations

Mass Percent Calculations

a. Calculate the grams of acetic acid (CH₃COOH) that would be present in one liter of your unknown solution. Use the mean value of the molarity from your sample.

________g CH₃COOH/L

b. Calculate the mass in grams of 1.000 liter of the acetic acid solution. The density of vinegar is 1.002 g/mL.

________g solution/L

c. Calculate the mass percent CH₃COOH in your sample.

________% by mass

Problems

1. A 27.86 mL sample of 0.1744 M HNO₃ is titrated with 29.04 mL of a KOH solution. What is the molarity of the KOH?

2. A 0.2977 g sample of pure H₂C₂O₄•2H₂O crystals is dissolved in water and titrated with 23.73 mL of a sodium hydroxide solution. What is the molarity of the NaOH?

3. A 0.3811 g sample of KOH will just neutralize what volume of 0.2000 M H₂SO₄?

4. A 0.8500 g sample of commercial lye ( a alkaline mixture that contains NaOH) , is dissolved in water and titrated with 35.00 mL of 0.5377 M HCl. What is the percent purity of the lye sample? (i.e., what is the mass percent of NaOH in the lye?)

5. A 30.0 mL aliquot of 0.300 M H₃PO₄ is mixed with 90.0 mL of 0.200 M KOH, and the mixture is evaporated. Will the salt that crystallizes out be K₃PO₄, K₂HPO₄, or KH₂PO₄? Show your calculations. (Hint: when a polyprotic acid such as H₃PO₄ is mixed with a base, the protons will react one at a time. For example, if the one mole of OH^-1 was mixed with one mole of H₃PO₄, it would produce one mole of H₂PO₄^-1 and one mole of water. The next proton on the H₂PO₄^-1 would start to react only when more OH^-1 was added, over and above the first mol.)