2. The student will gain experience using the buret (volumetric).
3. The student will have experience in the use of visual acid-base color indicators to determine the equivalence point of a titration.
4. The student will learn the analytical techniques of both forward and backward titration for the determination of a component.
5. The student will use the analytical balance to make sample weighing.
6. The student will be introduced to the quantitative calculations of analytical chemistry.
Skoog, West, and Holler: Fundamentals of Analytical Chemistry,
7th Ed., Saunders, 1996.
Fritz and Schenk: Quantitative Analytical Chemistry, 4th Ed., Allyn & Bacon, 1979, pages 548-553.
Fisher and Peters: Quantitative Chemical Analysis, 3rd Ed., Saunders, 1968, pages 65-66, 312-317.
J.M. Miller and D.V. Zahniser, Chemistry 44 (7) 28, 1971.
A. Analytical Balance
Purpose: Since the analytical balance is a fundamental measuring instrument for all types of quantitative work in analytical chemistry, a thorough knowledge of and familiarity with the construction and design, the theory of operation, and the inherent accuracy of a balance are important. In this experiment, the precision of weighings made with an analytical balance will be investigated and the magnitude of various common errors encountered in weighing will be evaluated.
2. Perform the preliminary zero adjustment.
3. Determine the weight of a clean, dry weighing bottle (without its lid) to with ± 0.1 mg.
4. Determine the weight of the weighing bottle lid separately to within ± 0.1 mg.
5. Determine the total weight of the weighing bottle plus the lid to within ± 0.1 mg. Compare this result with the sum of the separate weighings of the bottle and the lid to see how closely the weights agree.
6. Reweigh the weighing bottle without its lid to see how reproducible this weight is.
7. Roll the weighing bottle around in your hand: i.e., handle it and finger it and then reweigh it and compare the weight with the previous results. Next, wipe the weighing bottle clean with a dry, lint-free cloth or with laboratory tissue, reweigh it, and again compare the result to previous ones.
8. Hold the weighing bottle an inch from your mouth and breathe on it several times. Weigh it again and compare the result with earlier weights.
9. Place the weighing bottle in a drying oven (165o C) for two or three minutes. Remove the bottle from the oven with tongs and reweigh it immediately while it is still warm. Follow the change in its apparent weight for several minutes, recording the weight every 30 seconds.
10. Weigh the weighing bottle lid. Remove the lid from the balance and write your initials with a pencil on the ground-glass surface. Reweigh the lid and note the difference in weight.
2. List some possible errors in weighing and give an estimate of the size of each error.
3. Account for the results observed in step 5 of the
1. Hydrochloric Acid -- Approximately 0.1 M
Calculate the amount of 36% hydrochloric acid with a density of 1.2 g/ml that is needed to make 2 liters of 0.1 M hydrochloric acid. Check with the instructor. Add the volume of distilled water needed (2000 - volume of conc. acid in ml.) to your clean BROWN BOTTLE. Add the needed volume of conc. acid to the distilled water. Stopper. Shake well and let it stand over night. Use a liter graduated cylinder to measure the distilled water and a smaller graduated cylinder to add the concentrated acid. Measure the concentrated acid under the hood. Why? Label the bottle. (If your brown bottle holds 1900 ml. this is adequate.)
2. Sodium Hydroxide -- Approximately 0.1 M
One pound of sodium hydroxide pellets was dissolved in distilled water and the resulting solution occupied 900 ml. Calculate the molarity of this solution. Calculate the volume of this solution needed to prepare 2 liters of 0.1 M sodium hydroxide. Check with the instructor. Add the volume of distilled water needed (2000 - volume of conc. sodium hydroxide in ml.) to clean (distilled water rinsed) clear glass (or plastic) bottle. Add the needed volume of conc. base. Shake well. Let the solution sit overnight. Should you measure the conc. base in the hood? If so, why? If not, why not? Label the bottle.
C. DETERMINATION OF ACID BASE RATIO
To obtain the relative strength of your acid and your
base you need to determine the ratio of the two molarities. This is obtained
by placing 20 ml. of your acid in a 250-ml. erlenmeyer flask, adding about
100 ml. of distilled water, 3 drops of indicator [do the ratio to get three
results with phenolphthalein and then to get three results with the indicator
methyl-red-bromocresol green (MRBCG) and titrate with your base]. The phenolphthalein
endpoint is the first permanent pink color to persist for 10 seconds. The
methyl red-bromocresol green color does not fade so your want the first
color change. It is necessary to repeat the acid-base ratio until you get
at least 3 results to agree to within 3 parts-per-thousand relative deviation
with each indicator. Be patient. To get reproducible work in one of the
objectives of doing the acid base ratio.
D. STANDARDIZATION OF BASE WITH POTASSIUM ACID PHTHALATE
Weigh out samples (start with 4) of about 0.8 g. (you must know the weight to the tenth of the mg, but it may vary from 0.6 to 0.9 g) of dry potassium acid phthalate (KHP) in a dry 250 ml. erlenmeyer flask. Add 100 ml. (using your graduated cylinder) of distilled water. Add 3 drops of phenolphthalein indicator. Titrate to the first permanent pink color for 10 sec. Record the volume taken for the titration. Calculate the molarity of the base solution. The number of moles of KHP is the same as the number of moles of base in the volume of titrant used in the titration. Calculate your results. Prepare a spread sheet with your results and including the computer print-out with your report. If rejections or If new runs are needed, a new spread sheet is also needed.
E. DETERMINATION OF UNKNOWN ACID
Weigh out samples of approximately 0.7 g (to the nearest 0.1 mg. on the analytical balance) of unknown acid in a 250 ml. erlenmeyer flask. Add about 100 ml. of distilled water and 3 drops of indicator (phenolphthalein). Titrate to the first permanent pink color (for 10 sec.) and record the volume required for the titration. Calculate the percent KHP in the unknown acid. A spread sheet is also needed for this calculation.
F. STANDARDIZATION OF ACID
Calculate the molarity of the standard acid solution by using acid-base ratio. Remember that moles of acid = moles of base, and molarity x volume = moles. Therefore MaVa = MbVb. If you desired to standardize the acid chemically it is done by using a known purity sodium carbonate. This calculation will take less than 30 minutes. Ask for help, if needed.
G. DETERMINATION OF ANTACID CONTENT
Obtain six antacid tablets from the instructor. Weigh a single antacid tablet. Place several tablets in a dry mortar and grind the tablets to a powder. Desiccate the ground tablets over night. Weigh our about one-sixth of the powder and place it in a 250 ml. erlenmeyer flask. Add 100 ml of your standardized 0.1 M HCl solution with a buret and either add 3 drops of methyl red/bromo- cresol green indicator, or see the alternative procedure below. Did the solution turn red? If so, swirl and boil for one minute. If not, add 50 ml more acid. Did the solution turn pink? (or red?) If so, swirl and boil one minute. If the boiled solution is still pink after cooling back to room temperature, titrate is your standardized 0.1 M NaOH. NOTE: If a solid remains after boiling the sample, it may be necessary to use the alternate method below. See below for help! Calculate % Na2CO3.
%Na2CO3 = [ml. acid - (ml.
base x a/b ratio)](M acid)(1/2)(106)(100
weight in milligrams of powdered tablet
Run at least 4 tablets that agree to 10 ppt. Also calculate the grams and milligrams of HCl neutralized by 1 gram of antacid. Be prepared to share your results with the class by April 11.
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