CHM110H5 Lecture Notes - Equilibrium Constant, Titration, Sodium Hydroxide

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skykoala223

6 Dec 2013

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Lab 4 - identification of an unknown acid by ph titration. Purpose: to experimentally determine, through titration, the ka of an unknown acid in order to find out what the acid is. This will be determined by finding the ka for the following reaction: From there, it will be compared with the ka values of other acids. Experimental method: refer to pages 68-70 in course manual for chm110. In the beginning of the titration of the unknown acid, 0. 4 ml of naoh was added for the first reading instead of. 0. 25 ml, and proceeded as described in the manual from there. Additionally, in the titration of the unknown acid, the portion where naoh was supposed to be added drop by drop was skipped. Based on figure 1, when ph =7, 25. 9 ml of naoh was added. Ca = 0. 103 m: based figure 3, the equivalence point was reached at 24. 5ml.

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Could someone please help me with this chemistry problem? Thank you!

After performing a titration using 20 mL of unknown weak acid with 0.3830M NaOH, I obtained the following results:

0 mL added NaOH = pH 3.80

Final amount of added NaOH = 8.59 mL = pH 12.20

At Equivalence Point: Volume of NaOH = 4.835 mL; pH = approx 8.40

At the Halfway Point; Volume of NaOH = 2.418 mL: pH = 5.15

Ka = 7.08 x 10^-6

Initial [HA] = 0.0926M (determined by M1V1=M2V2 at the equivalence point)

Molar Mass = 1079.8 g/mol (determined by using original mass 1.9977 g/1.85 x 10^-3 moles

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Now I am asked the following questions in order to check my results for consistency;

Using your experimental Ka value, the concentration of your unknown acid solution and the concentration of the standardized NaOH solution, CALCULATE the pH of the following points on your titration curve:

1) The beginning of the titration before any any NaOH was added

2) 25% of the way to the equivalence point (25% of 4.835 mL is 1.21 mL)

3) 75% of the way to the equivalence point

4) The equivalence point (remember, this is found using Ka and concentrations of acid and base)

5) 20% of the way to the equivalence point (1.20 x 4.835= 5.80 mL)

I've found some info on the internet that says when in the buffer region, the moles of A- = moles of base added and that moles of HA = original moles of HA - moles of base added. Where I got confused was when the equilibrium expression was written as Ka = [H+] [mols Base Added/Volume] Ã· [mols Acid Left/Volume]. I don't know what volumes to use because I always think it makes more sense to use total volume of acid plus base mL's. But maybe you are supposed to use only the volume of each (acid/base) separately, in which case I think the volume of acid would be 20 mL. Maybe this is a totally wrong way to do it- I don't know! Please Help!!!!

I am unable to solve these questions due to a time constraint. I received help on part 1 of the unknown acid. Need help with the rest! I really appreciate it!

Lab: ACID-BASE TITRATION

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)

CHM110H5 Lecture Notes - Equilibrium Constant, Titration, Sodium Hydroxide

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