Question

I have a few questions from a lab experiment that my class has done for Acid-Base Titrations.

Objective:

To use titration to determine in Part 1, the concentration (molarity) of an unknown acid solution and in Part 2, the purity of a sample of KHP acid.

Background:

Titration is a volumetric technique used to determine the concentrations of solutions, molar masses of solids, and purity of samples. A titration inolves the addition of a titrant (a solution of known concentration) to an analyte ( a solution of unknown concentration), or vice versa, using a piece of glassware called a burette. The titration is carried out until it reaches an equivalence point (the exact point at which the reaction between the two solutions is complete). A chemical indicator is often used to aid in the identification of the equivalence point. An indicator changes color upon reaching the equivalence point of endpoint. Since we will be carrying out an acid-base titration, the indicator must change color upon reaching the endpoint at a specific pH. An example of such an indicator is phenolphthalein, which changes from colorless to pink near the equivalence point when pH approaches 7. It should be noted that the indicator selected is dependent upon a given titration, so that the observed color change is close to the ideal equivalence point. Once the titration is completed, we can use the volumes measured for each solution, as well as the concentration of the titrant, to determine the concentration, molar mass, or purity of the unknown analyte.

Procedure

Note: During lab, record all measurements and data in a clearly labeled table.

Part 1:

1. Obtain ~80mL of ~0.1M of base (sodium hydroxide, NaOH) into a labeled beaker.

Make sure to report the actual concentration from the bottle.

2. Obtain ~25mL of unknown acid into a second labeled beaker. Record whether your acid is monoprotic, diprotic, or triprotic.

3. Using a funnel, or carefully pouring, fill the burette with sodium hydroxide NaOH. Make sure a labeled waste beaker is under the buretter to catch any drippings. Rinse your burette by filling it with ~5-10mL of NaOH. Turn the knob to let ~2 mL out the tip into the waste beaker. Make sure your waste beaker is under the burette to catch any drips. Next, turn the Burette upside down and pour the remaining NaOH into your waste beaker while twisting the burette to coat the inside with NaOH.

4. Fill your burette with sodium hydroxide. Turn the knob to fill the tip of the burette with NaOH. Record the initial volume to the correct number of significant figures.

5. Using a volumentric pipette, transfer 10.00 mL of your unknown acid into the Erlenmeyer flask.

6. Add 3 drops of pH indicator and magnetic stirrer to the Erlenmeyer.

7. Under the burette, place the Erlenmeyer on a hot plate and turn on the stir knob. Heat should be off.

8. Titrate the unknown acid until the solution turns pink and remains pink.

9. Record the final volume on the burette to the correct number of significant figures.

10. Repeat (4) to (7). Pre-plan the whole lab and who will do which task.

MY RESULTS:

PART 1

The actual concentration of NAOH is 0.5939, 80 mL

Initial	Final
24.9 mL	44.5mL
26.1 mL	45.57mL

PART 2

Impure KHP (SAMPLE B )

Only Paper: 0.280 g

Impure KHP Mass: 0.526 g

Initial volume: 22.30 mL + water +26.10 mL

Final Volume: 26.80 mL

QUESTIONS:

1. Calculate the concentration (molarity) of your unknown acid.

2. Calculate how many grams in your impure KHP sample was indeed KHP.

3. Calculate the percent (by mass) KHP in the sample.

Part 2:

1. What would happen to the final result( state what the final result is ) if the tip of the burette was not filled when reading the initial volume?

2. What would happen to the final answer if you used Ca(OH)2 instead of NaOH and DID NOT know it?

3. What would happen to the final result if the Erlenmeyer flask was contaminated with other types of acid?

4. Recalculate the concentration of your unknown acid if it was given as a diprotic acid.

5. Draw a molecular level picture (showing Na+, H+, OH-, H2O, etc...) of what is in your beaker before, at half point, at equivalent point, and at end point of the titration.

Part 3:

1. What would have happened to the final result if the impurity in the sample was also acidic?

2. What would have happend to your final result if you used ~25mL of water to dissolve the solid acid instead of 20mL?

Summary data:

Create a clearly labeled table of all your final data results for each run, and athe average of multiple runs.

Part 1: Unknown #, Vol base, M base, Vol acid, M acid.

Part 2: Sample #, % KHP (purity)

Answer 1

Calculation for unknown acid

Given

NaOH concentration 0.1 M (0.1 moles/litre), Molecular weight of NaOH 40 g/mol

That means 4 grams of NaOH is there in 1000 mL of Water

Buret reading 19.6 mL (used volume of 0.1 M NaOH in titration)

To determine the number of grams of NaOH used, apply cross multplication

Number of moles of NaOH

At the end point number of moles of NaOH is equal to the number of moles of unknown acid.

So, number of moles of unknown acid is 0.002 mols

And the given volume of unknown acid is 25 mL (0.025 litres)

Concentration (M) of unknown acid

Calculation for KHP

KHP Potassium Hydrogen Pthalic acid. It is a mono protic acid (capable of donating only one proton). So, one equivalent of NaOH is for one equivalent of KHP at the end point

Molecular weight of KHP 204.22 grams/mol. That means 40 grams of NaOH takes up 204.22 grams of KHP

Reading of 0.1 M NaOH is 26.80 mL (given above as final volume)

To know the number of grams of KHP for 0.107 grams of NaOH

That means 0.546 grams of KHP (pure) reacted in titration

Given impure mass of KHP = 0.526 grams (is it a typo ?)

For purity calculations

%