Standardization of Acids and Bases

Description

Solutions of known concentration are prepared by dissolving measured masses of standard acids in distilled water. The concentrations of unknown solutions of sodium hydroxide are determined by titration.

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• An acid solution reacts with a base solution in a "neutralization" reaction.
• The pertinent chemical reactions in this experiment are:

  KHP + NaOH --> KNaP + H₂O

  

  H₂C₂O₄ + 2NaOH --> 2H₂O + Na₂C₂O₄

  H₂SO₄ + 2NaOH --> 2H₂O + Na₂SO₄

  H⁺ + OH^- --> H₂O

• The volume of solution required for such a reaction to go to completion may be measured using a buret.
• A color change in an added chemical indicator signals an end point. It is selected to indicate that neutralization has taken place.
• This end point may not be the same as the stoichiometric or equivalence point; given an appropriate selection of indicator, the difference in volume between these two points should be negligible.
• In this experiment, a solid acid is titrated with sodium hydroxide to the faint pink phenolphthalein color.
• Once the concentration of the NaOH is known, it may be used to calculate the concentration of other acids.
• The procedure of determining precisely the concentration of an acid or base solution is called standardization -- using a known concentration to determine an unknown concentration.
• The product of molarity (mol/L) times volume (L) gives moles.
• For a solid acid, the moles of H⁺ is given by:

  moles of H⁺ = mass (g) x (mol H⁺/mol acid) x ( 1 (mol)/ molar mass (g))

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Sodium hydroxide damages tissue and causes blindness. Oxalic acid, phenolphthalein, and potassium hydrogen phthalate are toxic.

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Wear eye protection. Do not ingest chemicals.

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Solid Acid

1. Follow the "Titration Skills Checklist." (below)
2. Place approximately 3 g of solid oxalic acid in a weighing container.
3. Clean three 250-mL Erlenmeyer flasks and label them 1 through 3.
4. Place a carefully weighed sample of roughly 1.0 g into each flask by using the following method:
5. Weigh the container and solid acid as accurately as your balance will allow.
6. Record this weight in the Data Table.
7. Remove 1 gram from the balance reading.
8. Transfer solid to the flask carefully by tapping and checking the weight to see if it is close. Initially the vial will be too heavy for the weights on the balance.
9. Tap and reweigh until about 1 g is transferred to flask 1. (It is not necessary to have exactly one gram in each flask but it is necessary to know accurately how much you have in each flask.) The balance will be swinging slowly if a triple beam balance is used.
10. Reweigh the container accurately.
11. Record in the Data Table.
12. Repeat for flasks 2 and 3.
13. Add 30 mL of distilled water. Swirl. Do not be concerned if the solid acid does not dissolve.
14. 2 to 3 drops of phenolphthalein to each flask.
15. Rinse the buret with three 5-mL portions of the NaOH solution which you have in the large flask.
16. Fill the buret to a convenient volume between the 0-mL and 1-mL mark. (Some workers use a funnel in the buret; others do not.)
17. Make sure the tip is full. If not, twist the stopcock very rapidly to remove air bubbles.
18. Use a 3" x 5" card with a thick black stripe drawn on it. Place this card behind and just below the meniscus. A reflection of the stripe appears on the liquid surface and makes the meniscus more visible. Record the initial buret reading in the Data Table.
19. Place the titration flask under the buret.
20. Open the stopcock to add base to that the titration can begin. Notice how the stopcock is held.
21. It is essential to swirl the solution continuously during the titration process.
22. Periodically stop the flow and rinse the walls of the flask with distilled water.
23. Continue adding sodium hydroxide solution more and more slowly until one drop of titrant causes a pink color that does not disappear even with agitation.
24. Record the volume in the buret.
25. To check, add one more drop to be sure the solution was pink.
26. Record the final reading in the Data Table.
27. Repeat for Flasks 2 and 3.
28. Clean the flasks.

Acid Solutions

1. Label two flasks 1 and 2.
2. Pipet 25 mL of the sulfuric acid solution into each flask.
3. Add 25 mL of distilled water and 2 to 3 drops of phenolphthalein to each flask.
4. See the detailed titrating instructions for the solid acid.
5. Record the initial buret reading in the Data Table.
6. Titrate the sulfuric acid solution in Flask 1 to the lightest pink that your eye can detect and that will not go away with stirring.
7. Record the final buret reading in Table 2.
8. Follow the recording and titrating steps for the remaining sulfuric acid solution.
9. Complete the tables, show sample calculations, and answer the follow up questions.

Titration Skills Checklist

______ Buret clean before use

______ Air excluded from tip

______ Starting volume correctly recorded

______ Unknown solution correctly measured

______ Indicator added

______ Stopcock held correctly

______ Endpoint satisfactory

______ Final volume correctly recorded

______ Buret rinsed after use

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Name ___________________________ Class ________

Teacher__________________________

DoChem 102 Standardization of Acids and Bases

Titration Skills Checklist

______ Buret clean before use

______ Air excluded from tip

______ Starting volume correctly recorded

______ Unknown solution correctly measured

______ Indicator added

______ Stopcock held correctly

______ Endpoint satisfactory

______ Final volume correctly recorded

______ Buret rinsed after use

_______________________________________________

________________________________________________

Closure Questions:

1. List sources of error for this experiment.
2. A 0.300 M solution of NaOH cannot be made up by weighing solid NaOH and dissolving in the appropriate amount of water. Suggest reasons why.
3. Why is phenolphthalein a suitable indicator for this experiment?
4. Identify other indicators that might be used.
5. Find the molarity of a solution containing 2.00 g of oxalic acid dissolved in 100.0 mL of H₂O.
6. If 25.00 mL of NaOH were used to neutralize 32.0 mL of the acid from the solution described above, what is the molarity of the NaOH?

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Name ___________________________ Class ________

Teacher__________________________

DoChem 102 Standardization of Acids and Bases

Watch the movie and record the buret readings.

1. Record the initial reading of the buret on the movie. _____
2. Record the final reading of the buret on the movie. _____

Use the data below to calculate the molarity of the NaOH solution.

4. Calculate the molarity of the sulfuric acid solution titrated below.

Closure Questions:

1. List sources of error for this experiment.
2. A 0.300 M solution of NaOH cannot be made up by weighing solid NaOH and dissolving in the appropriate amount of water. Suggest reasons why.
3. Why is phenolphthalein a suitable indicator for this experiment?
4. Identify other indicators that might be used.
5. Find the molarity of a solution containing 2.00 g of oxalic acid dissolved in 100.0 mL of H₂O.
6. If 25.00 mL of NaOH were used to neutralize 32.0 mL of the acid from the solution described above, what is the molarity of the NaOH?

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Purpose

• Prepare and standardize a base solution
• Use titration techniques.
• Standardize an acid solution.

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(per 10 students working in pairs)

• 2 L 0.5 M sodium hydroxide solution (20 g NaOH/ L)
• 2 L 0.250 M sulfuric acid (14.0 mL 17.8 M H₂SO₄ (95%)/ L solution)
• 5 small weighing container
• 5 balance
• 5 500-mL flask with stopper
• 30 g of potassium hydrogen phthalate (KHP) or oxalic acid
• 15 Erlenmeyer flask, 125- or 250-mL
• 5 buret, 25- or 50-mL
• 5 wash bottle with distilled water
• 5 10-mL dropper bottle of phenolphthalein (1 g/(50 mL ethanol + 50 mL H₂O)) or other suitable indicator
• 5 50-mL graduated cylinder
• 5 25-mL volumetric pipet and pipet bulb

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• Have students use the "Titration Skills Checklist."
• The weight of the solid acid that should be used in each flask will vary with different acids. For 40 mL of 0.5 M NaOH, the mass of solid acid is:

  mass = 0.020 x molar mass acid x (mol acid/ mol H⁺)

• The larger the molar mass of the standard solid acid, and the fewer titratable protons per mole, the larger will be the required mass. Mass error is the largest error in the experiment when using centigram balances.
• The teacher should prepare the sodium hydroxide solution for the students. Place 40 mL of 6 M NaOH solution in a 500-mL flask and add enough distilled water to make 500 mL of solution. This will give you a solution approximately 0.5 M. An alternate procedure is to dissolve 20 g solid NaOH per liter of solution to be prepared. Use caution; the glass container holds a liquid that can cause great damage to eyes. Shake this solution well with the stopper in place.
• Clean the burets yourself after the lab is finished; less breakage will occur.
• Rinsing with three 5-mL portions of the solution that is to be put in the buret is sufficient to clean the buret if the buret was clean before the first class used them.
• Have the students save the standardized NaOH solution if you plan to do more titration experiments.
• Pipeting volumes of H₂SO₄ is much more accurate than using a graduated cylinder.
• A standard acid such as 0.300 M HCl (43.1 mL of 11.6 M (36%) HCl/ L H₂O) could be used instead of a solid acid. This will usually save time. The instructor should do the experiment in order to standardize this acid for the students use or the student can be told it is 0.300 M even if it is only 0.3 M. Be sure to have students pipet for accurate volumes.

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Teacher preparation: 30 min

Class time: 2 periods

• 40-50 min first period
• 30 min second period (acid titration during second period)

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Wear eye protection. Provide a working eye wash. An emergency shower should be available. Do not ingest chemicals.

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The materials used in this experiment may be disposed of safely at the sink after neutralization.

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Closure Questions:

1. List sources of error for this experiment.
2. A 0.300 M solution of NaOH cannot be made up by weighing solid NaOH and dissolving in the appropriate amount of water. Suggest reasons why.
3. Why is phenolphthalein a suitable indicator for this experiment?
4. Identify other indicators that might be used.
5. Find the molarity of a solution containing 2.00 g of oxalic acid dissolved in 100.0 mL of H₂O.
6. If 25.00 mL of NaOH were used to neutralize 32.0 mL of the acid from the solution described above, what is the molarity of the NaOH?

Answers to Closure Questions:

1. Weighing errors supersede all other errors when a centigram balance is used. Weighing contributes 2-3% error at the outset when oxalic acid is used as the standard. The weighing error will be smaller with potassium hydrogen phthalate.
2. NaOH absorbs both water and carbon dioxide. It is too difficult to keep chemically pure to be used as a standard.
3. Phenolphthalein changes color in a solution where the [OH^-] is slightly higher than the [H⁺]. Square braces, [xx], are used to represent molar concentrations: [H⁺] is read as "The molar concentration of H⁺."
4. Bromthymol blue, methyl red, and methyl orange may be used as indicators for the sulfuric acid titration. Phenolphthalein is the best choice for the oxalic acid titration.
5. Molarity = (mol/L)

   =2.00g H₂C₂O₄ x (1 mol H₂C₂O₄ / 90.00 g H₂C₂O₄) x (1/ 100 mL) x (1000 mL/ L)

   = 0.222 M.

6. 32 mL x (1 L/1000 mL) x (0.222 mol H₂C₂O₄/ 1 L)

   x (2 mol H⁺ / 1 mol H₂C₂O₄ ) x (1 mol NaOH/1 mol H⁺)

   x (1/ 25 mL) x (1000 mL/ 1 L)

   = 0.568 mol.

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1. 19.83
2. 0.33
3. and 4. see sample data for the calculated values.

   See closure question answers.

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• acid
• base
• standard acid
• titration
• neutralize
• color change
• end point
• concentration
• molarity

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