A Demonstration of the Titration Process in a Lab

Part A and B demonstrate the variations between results in phenolphthalein and methyl orange indicator. Based on the titrations in Part A and Part B (sections in order to figure out how to investigate the unknown amount of an acid or base in a solution) the unknown amount in Part C is 89.33% basic.

Introduction

Titration is the reaction of a known amount of a substance with an unknown amount of another substance in order to determine its concentration. For this process to be of value to the user the chemical reaction must be very well defined and there must be some observable change accompanying the reaction so that the user is able to monitor the progress of the reaction. Sometimes this observable change involves a color change, sometimes a physical change, sometimes a pH change, and sometimes other factors can be monitored. In today’s chemical research and industrial environment, this process is often automated.

An “indicator” is often used in titrations to monitor the progress of the reaction. They are especially common in acid/base titrations where they indicate changes in pH over specific pH ranges is by a change in color. The indicator does not participate in the ration, and is usually only added to the solution in very small amounts so that its influence on the reaction is minimal. You might think of the indicator as an outside observer to the reaction.

Classical titration involves the use of a buret. In order for a titration to be useful, the concentration of one of the two reactants must be known precisely. The titration will involve the addition of an unknown concentration of hydrochloric acid to a known amount of sodium carbonate.

Materials

• 50 mL Cylinder
• 3 125 mL Flasks
• Balance
• Funnel
• Buret

• o Stand

• o Clamp

• Wash bottle
• dH¬2O
• Na2CO3
• HCl
• Phenolphthalein Indicator
• Methyl Orange Indicator
• Gloves
• Weigh Boats
• Measure Shovel
• Protective Glasses/Goggles
• Unknown mixed sample of NaCl & Na2CO3
• Stirring Plate

Methods

Part A

Acquire three 125 mL flasks, weigh out close to 0.2 grams of sodium carbonate. Make sure you record the weight of sodium carbonate as exactly as your balance allows so that you can find the unknown concentration of HCl more accurately.

Add about 20 mL of water to each of the flasks.

Add two or three drops of phenolphthalein indicator to each of the flasks.

Titrate your “known” solutions of Na2CO3 by carefully adding the unknown solution of HCl from the buret. Note that phenolphthalein changes from pink to clear just after the first equivalence point of the reaction has been reached; that is, when enough acid has been added so that the moles of acid and base in your solution are equivalent.

Calculate the concentration of HCl in the unknown solution from the volume of unknown solution required to titrate you three known solutions. Report the average concentration from your three titrations with a standard deviation.

Part B

Repeat Part A, this time adding two or three drops of methyl orange indicator instead of phenolphthalein in step 3. Note that methyl orange changes from a yellow color to a pinkish “sunset” color at the second equivalence point of the reaction.

Part C

Use what you have learned to identify the concentration of Na2CO3 in an unknown mixed sample of NaCl and Na2CO3. Use the solution of HCl (uknown 1) that you determined the concentraton of, as a standard, to titrate the Na2CO3. Report the concentration in weight %.

Results

Part A

Flask A Flask B Flask C

Initial Mass (g) of Na2CO3 1.79 1.92 1.94

Initial Volume (mL) of HCl 4.2 21.1 18.1

Final Volume (mL) of HCl 21.1 39.3 36.5

Volume Used of HCl 16.9 18.2 18.4

Molarity of HCl 0.117 0.104 0.103

Average: 0.108 M HCl

Part B

Flask A Flask B Flask C

Initial Mass (g) of Na2CO3 .2011 0.2008 .2011

Initial Volume (mL) of HCl 0 6.3 0.1

Final Volume (mL) of HCl 36.1 42.3 40.4

Volume Used of HCl 36.1 36.0 40.3

Molarity of HCl 0.105 0.105 0.094

Average: 0.101 M of HCl

Part C

Flask A Flask B Flask C

Initial Volume (mL) 7.0 23.4 9.3

Final Volume (mL) of HCl 23.4 40.5 26.3

Volume (mL) Used of HCl 16.4 17.1 17.0

Flask A

0.1 mol HCl x 0.0164 L HCl x 1 mol Na2CO3 x 105.96g Na2CO3 x __________ = 0.868g = 87%

1 L 1 mol HCl 1 mol Na2CO3 0.2g Na2CO3

Flask B

0.1 mol HCl x 0.0171 L HCl x 1 mol Na2CO3 x 105.96g Na2CO3 x __________ = 0.905g = 91% 1 L 1 mol HCl 1 mol Na2CO3 0.2g Na2CO3

Flask C

0.1 mol HCl x 0.017 L HCl x 1 mol Na2CO3 x 105.96g Na2CO3 x ___________ = 0.900g = 90%

1 L 1 mol HCl 1 mol Na2CO3 0.2g Na2CO3

Average: 89.33%

Discussion

In this experiment, the indicators are used to determine when a solution has become neutralized or reached a point where it is slightly more acidic (pending human error). The phenolphthalein tells us when the solution is basic and when the color of the indicator changes, then that indicates the solution is neutral. The methyl orange indicator also implies that the solution is basic as well, but when the color of the indicator has disappeared then the solution is neutral. The bubbles are from a chemical reaction between the HCl and the Na2CO3 turning into NaCl, water, and Carbon dioxide. The bubbles are the carbon dioxide molecules being released into the environment. It is important to perform each titration more than once in order to have the best results. The variation in the results are due to the human eye viewing the change in indicator color therefore, the person performing the experiment will see the change in color at different times and at different volumes. The results have been repeated more than once so they are close to being accurate as they can be even with human error. Given that there is human error, they are off by a few mL.