Unit 13 Acids and Bases

Transcription

1 Unit 13 Acids/Bases Acids can be simply defined as compounds that can produce H + ions and generally have an "H" as the first element in the formula (e.g. HCl, H2SO4, HNO3, etc.). Bases are simply defined as compounds that produce OH - ions or have an "OH" group at the end of the formula (e.g. NaOH, KOH, Ca(OH)2, etc.) Acids react with bases to form water (H + + OH - ---> H2O). More complex definitions exist but will not be needed for the next few experiments involving acids and bases. Since most acid/base reactions occur in the liquid state, we will need to introduce the concept of "titration" and solution concentration in order to perform quantitative determinations. Solution concentration can be expressed in several ways. A w/w percent is commonly used in commercial practice. However, chemists find that other methods of expressing concentration are more convenient. Molarity is one of the most commonly used. It measures concentration as the moles of solute (minor component) per liter of solution. It is abbreviated using a capital "M". Thus 16M means 16 moles of solute per 1 liter of solution. A similar expression for concentration is normality. It is the moles of H + or OH - per liter of solution and is abbreviated with a capital "N". Thus 0.555N means that there are moles of H + or OH - per liter of solution. The later is convenient for chemists working with acids and bases since it eliminates the need to know the source of the acid or base, whether it be a monoprotic, diprotic, or triprotic acid, or monohydroxy, dihydroxy, or trihydroxy base. A titration involves the use of a burette, a graduated cylinder that allow filling at the top and dispensing at the bottom. The graduations begin with "0" at the top and increase downward which allow the measure of the solution delivered to another container rather than the amount of liquid actually in the burette. Most burettes are divided into 0.1 ml divisions. The diagram on the next page will help you read between these divisions. Remember the volume increases as we go downward. When we perform a titration we fill the burette with a titrating solution and add it drop wise to another solution in an Erlenmeyer flask. A burette holder is used to hold the burette so that two hands can be used to operate the stopcock. We will place a NaOH solution in the burette, measure the initial volume, add it drop wise to a flask

2 Unit 13 Acids and Bases containing an acid until all the acid has been reacted. Then, by knowing the amount of acid in the flask, we will be able to determine the concentration of the titrating solution in the burette. How are we going to know when all the acid has reacted with the base? We cannot see the water forming. We can do this using an acid/base indicator. We will use phenolphthalein (1% in a 50/50 mixture of alcohol and water). When acid is present it is colorless. When all the acid is gone and a base remains it is pink to red in color. We will add the base to the flask containing the acid while swirling or 128

3 Unit 13 Acids/Bases mixing until a persistent pink color remains for at most 30 seconds. Acids and bases often are hydroscopic. This means they absorb moisture from the air. Bases also absorb carbon dioxide from the air. When carbon dioxide from the air dissolves in water it forms carbonic acid that can neutralize bases. For these reasons, NaOH solutions are considered unstable unless the air above them can be dried and void of carbon dioxide. In industrial labs, Drierite (or other suitable desiccant) and Ascarite (carbon dioxide absorbent) are used to treat air drawn into a vessel containing a base as the base is used in titration. Such protected solutions are stable and will not change their concentration. In our lab, we will make an NaOH solution, standardize it, and keep it in a sealed container and use it before it can be significantly contaminated with atmospheric moisture and carbon dioxide. Unstable solutions usually need to be standardized. Standardization is the use of a known material to determine the concentration of an unknown material. We will use potassium hydrogen phthalate abbreviated KHP. This is not a chemical formula and there is no phosphorus in the compound. It is just a convenient way of writing the name of the acid. Its molecular mass is g/mole. KHP is the national standard for standardizing NaOH solutions. The reason for this is that it not hydroscopic. Its mass is stable even in moist air. By knowing the mass of KHP used, we will be able to calculate the moles of KHP, the moles of NaOH reacted to reach the phenolphthalein end point using stoichiometry, and then the Molarity or normality of the NaOH solution by dividing the moles of NaOH by the liters of NaOH used in the titration. Procedure 1. Prepare about 1 liter of a 0.1 N of NaOH by diluting a stock solution from the lab supply bench. Use the formula: 0.1 N NaOH needed x 1000 ml= ml stock soln needed? N NaOH stock soln 129

4 Unit 13 Acids and Bases Add this amount of the stock solution to a plastic 1 liter bottle (or equivalent) along with enough water to make the total volume 500 ml. Shake to mix and seal when not in use. Do not shake the solution again. This solution will need to be stored in your drawer since it will be used in the next several experiments. 2. Fill the burette with about 5 ml of the NaOH titrating solution. Turn the burette almost horizontal and rotate to rinse the burette with the solution. Discharge the solution through the tip to rinse the tip and discard the remaining solution. Refill the burette with the NaOH solution. This rinsing ensures that the solution in the burette will not be contaminated. Read and record the initial reading of the burette, which should be close to zero, to the second (hundredths) decimal place using the diagram in the discussion section above. 3. Measure about 0.1 g of the KHP into a medium sized Erlenmeyer flask. Record the exact mass on your laboratory report. Add about 10 ml of water to dissolve the KHP. Add a few drops of phenolphthalein indicator from the dropper bottles provided. 4. Add the NaOH solution drop wise while swirling the flask until a persistent pink color remains in the flask for at least 30 seconds. An end point should be reached with the addition of no more than one drop. If you overshoot the end point, you will need to repeat the titration. The second time you should have an idea of how much base will be needed and you can add the base rapidly until you approach the end point. At that point you will need to add the base one drop at a time. 5. When the end point has been reached record your final burette reading on 130

5 Unit 13 Acids/Bases your lab report. The volume titrated is the difference of the two readings. 6. Repeat the titration until you have two readings that give similar NaOH concentrations. Report the average of your best 2 or 3 concentrations. Keep your NaOH solution for the next two units. 131

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7 Prelab Exercises Unit 13 Unit 13 Acids/Bases Name Section Date 1. Why must we standardize the NaOH solution? 2. Why is KHP used to standardize the NaOH solution? 3. Why must we use an indicator in the titration? 4. What is the difference between 0.1 N and 0.1 M NaOH? 5. If we titrate 5.34 ml of the NaOH with about g of the KHP, what is the concentration of the NaOH solution? 133

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9 Unit 13 Acids/Bases Lab Report for Unit 13 Name Section Date Trial 1 Trial 2 Trial 3 mass of KHP used moles of KHP used moles of NaOH titrated final burette reading initial burette reading volume of NaOH titrated Normality of NaOH 135

10 Unit 13 Acids and Bases Average Normality: 136

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