Analysis of Commerical Antacids Containing Calcium Carbonate Prelab (Week 1)

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1 Analysis of Commerical Antacids Containing Calcium Carbonate Prelab (Week 1) Name Total /10 SHOW ALL WORK NO WORK = NO CREDIT 1. What is the purpose of this experiment? 2. Show the calculation for determining the mass of sodium hydroxide pellets required to prepare ml of a 0.10 M sodium hydroxide solution. You will need this value to make your sodium hydroxide solution in Part I. 3. You are going to standardize your sodium hydroxide by titrating with potassium hydrogen phthalate. As an example, you dissolve g of potassium hydrogen phthalate, otherwise known as KHP (KHC 8 H 4 O 4 ) in water in a ml Erlenmeyer flask and then add phenolphthalein indicator. You then titrate with your sodium hydroxide solution, which is in a buret, and you determine that the equivalence point is at ml of your sodium hydroxide solution. Determine the molarity of your sodium hydroxide solution. What color will the solution of potassium hydrogen phthalate turn to determine when the equivalence point has been reached? 1

2 Analysis of Commercial Antacids Containing Calcium Carbonate Prelab (Week 2) Name Total /10 SHOW ALL WORK NO WORK = NO CREDIT 1. You weigh out an antacid tablet and determine the mass to be g. After taking a g sample of the antacid tablet, you dissolve it in ml of a M solution of hydrochloric acid in a ml Erlenmeyer flask. After heating the solution to get rid of carbon dioxide, you titrate the leftover hydrochloric acid in the solution in the Erlenmeyer flask with a M solution of sodium hydroxide and determine that it takes ml of the sodium hydroxide solution to react with the leftover hydrochloric acid in solution. Determine the milligrams of calcium carbonate in the sample of antacid tablet and then determine the milligrams of calcium carbonate in the entire antacid tablet. 2. What indicator will be used for this titration reaction? What color change will occur? 2

3 Analysis of Commercial Antacids Containing Calcium Carbonate In this experiment, you will study acid-base reactions. You will determine the correct concentration (molarity) of a solution you make to be used for the ENTIRE experiment. This solution in turn, will be the key in determining the amount of calcium carbonate in an antacid tablet. Introduction In this experiment, you will analyze a commercial antacid for the amount of calcium carbonate present. The stomach contains among other things, hydrochloric acid. This is the acid that causes the gastric juices in the stomach to digest food at a ph of around 2 or 3. Many people believe that heartburn is caused by an overabundance of stomach acid, but this is not true. The stomach acid refluxing up through the esophageal sphincter, the valve separating the esophagus from the stomach causes heartburn. The esophagus does not have a mucus lining like the stomach. This lack of mucus causes a burning sensation due to the irritation by acid in the esophagus. Antacids give relief not by neutralizing the stomach acid, but by raising the ph to about 3 or 4. Most commercial acids contain one or more of the following neutralizing agents: calcium carbonate, sodium bicarbonate or salts of magnesium or aluminum. Antacid tablets also contain other ingredients to perform such tasks as holding the tablet together. Before you begin this experiment, there are many concepts and terms you will have to become familiar with. The first of these terms is molarity which is the measure of the concentration of a solution and has the symbol M. M = molarity = moles of solute litres of solution The solute is the material that is the dissolved in the solvent. In this case, our solvent will be water. You will be asked to make 250 ml of a 0.10 M solution of sodium hydroxide (NaOH). You can use this definition of molarity to determine what mass of NaOH must be weighed out to make 250 ml of a 0.10 M solution. The concept that you will become most familiar with in this experiment is that of acid-base chemistry. An acid, according to the Arrhenius definition, is a substance that contains hydrogen and releases hydrogen ions (H + ) while a base is a compound which releases hydroxide ions (OH - ) when they are dissolved in water. The result of an acid and a base reacting together gives the products of a salt and water. For example, when reacting HCl and NaOH together, the following reaction occurs: HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (l) acid base salt water This reaction is called a neutralization reaction. The stoichiometry of this reaction can change depending on the number of hydrogen or hydroxide ions that are released in the reaction. In the example above there is only one hydrogen ion and one hydroxide ion released in the reaction. There are acids as well as bases that release more than one hydrogen or hydroxide ions. The number of hydrogen ions that can be released into the solution by an acid is referred to as the acids proticity. An acid which can release one H + ion into solution is termed monoprotic and has a proticity of one. Examples of monoprotic acids are HCl, HC 2 H 3 O 2 (acetic acid) and HNO 3. An acid which can release two H + ions into solution is termed diprotic and has a proticity of two. Examples of diprotic acids are H 2 SO 4, H 2 C 2 O 4 2 H 2 O (oxalic acid) and H 2 CO 3. An acid which can release three H + ions into solution is termed triprotic and has a proticity of three. Examples of triprotic acids are H 3 PO 4 and H 3 C 6 H 5 O 7 (citric acid). 3

4 Not all substances are pure and stable enough to be weighed out directly and obtain a solution of accurately known concentration. Such is the case with sodium hydroxide, which rapidly absorbs water from the air. The method that will be used to standardize the sodium hydroxide, in order to determine its exact concentration is an acid-base titration. The acid used will be potassium hydrogen phthalate (KHC 8 H 4 O 4 ) also referred to as KHP. It is called a primary standard, meaning it is pure and stable enough to be weighed out directly in order to obtain a solution of accurately known concentration. In order to standardize your sodium hydroxide solution, you must understand the concept of titrating. A titration is defined as a procedure for the quantitative analysis of a substance by means of an essentially complete reaction in solution with a reagent of known concentration. The reagent of known concentration in our case is KHP. A buret will be used to perform the titration. The buret holds one of the reactants called the titrant and adds this into a reaction vessel which contains the other reactant. The titrant in this experiment will be the sodium hydroxide solution that you make. The KHP solution will be in the reaction vessel and will have a few drops of an indicator added to it at the start of the titration. An indicator is a dye that changes color when the reaction used for analysis is complete. The indicator you will be using in this experiment will be phenolphthalein, this indicator stays colorless until the reaction is complete and then it turns a faint pink. At some point in your titration, you will have added the same number of moles of OH - ions as there are moles of H + ions in the reaction vessel. In other words, there will be equal number of OH - and H + ions in the solution. This point in the reaction is called the equivalence point or endpoint. At this point, the indicator will change color when the slightest amount of titrant is added. It is at this point you want to stop and record this final volume. In Part I of this experiment, you will prepare your 250 ml of 0.10 M sodium hydroxide solution. Next, you will standardize this solution using it as a titrant with potassium hydrogen phthalate or KHP. The reaction that will occur is: KHC 8 H 4 O 4 (aq) + NaOH (aq) KNaC 8 H 4 O 4 (aq) + H 2 O (l) At the completion of the reaction, the moles of NaOH which you added from the buret can be determined from the grams of KHP you weighed out which you can convert to moles and then from the stoichiometry of the reaction. The molarity of NaOH can then be calculated by using the moles of NaOH which you calculate and the volume of NaOH that was added from the buret. In Part II of the experiment, you will perform what is called a back titration. You will dissolve the antacid tablet in a known amount of hydrochloric acid having a molarity of 0.1M. You will then titrate the solution with NaOH to determine how much hydrochloric acid was leftover. The difference between the moles of hydrochloric acid present before adding the antacid tablet and the moles of hydrochloric acid neutralized by the sodium hydroxide will be the amount of moles of hydrochloric acid neutralized by the antacid tablet. You can then use the stoichiometry to determine the number of moles and the mass of calcium carbonate that is present in one tablet of the antacid. In Part II there are two reactions you need to consider. The first reaction is the reaction between the hydrochloric acid and the calcium carbonate in the antacid tablet: 2HCl (aq) + CaCO 3 (s) CO 2 (g) + CaCl 2 (aq) + H 2 O (l) The second reaction is the reaction between the leftover hydrochloric acid and the sodium hydroxide: NaOH (aq) + HCl (aq) NaCl (aq) + H 2 O (l) 4

5 You can then compare your results with the amount of calcium carbonate per tablet stated on the list of ingredients for the commercial antacid you analyzed. This amount is generally around 500. mg/tablet. To calculate the mass of calcium carbonate in the entire tablet, you need the mass of your sample of tablet, the mass of the calcium carbonate in that sample of tablet and the mass of the entire tablet. You will set up the data according to the following equation and solve for the mass of calcium carbonate in the tablet. mass (mg) of CaCO 3 in sampleof tablet mass (mg) of CaCO 3 in tablet mass (g) of sample of tablet mass (g) of tablet Procedure Part I Preparing NaOH Solution 1. Rinse a 400 ml beaker with distilled water and dry it. Using a top loader balance, weigh out the approximate mass of NaOH that you need (you calculated this in your prelab) into the clean 400 ml beaker. Be careful when working with the NaOH it can cause severe burns, so NEVER touch the NaOH and clean any spills immediately. 2. Add enough distilled water into the beaker to make 250 ml of solution. Mix carefully with a stirring rod to dissolve all of the pellets. When all of the pellets are dissolved, carefully pour this solution into a clean polyethylene bottle. Standardizing NaOH Solution Data that you should have written down for EACH titration for Part I are: Titration Reaction Mass of KHP Moles of KHP Moles of NaOH used Final Buret Reading Initial Buret Reading Volume of NaOH used Molarity of NaOH solution Average Molarity of NaOH solution (For good titrations only.) Mass of antacid tablet. You should have sample calculations of just ONE titration shown in your laboratory notebook. 5

6 1. Clean a 250 ml Erlenmeyer flask and rinse it with distilled water. Using the analytical balance, weigh out between and g of KHP. Pour this into the Erlemeyer flask. Be careful not to spill any of the acid, try to get all of it into the flask. 2. Add 100 ml of water to the flask and swirl to dissolve all of the sample completely. Add 2-3 drops of phenolpthalein indicator. Most errors occur because people forget to add the indicator. 3. You must make sure that your buret is free of contamination since this can seriously affect your data. First clean and rinse the buret properly. When you are ready to begin the titration, rinse the buret with three 5 ml sample of the titrant (in this case, the NaOH solution). Let this solution drain out the bottom by opening the stopcock. Make sure the whole interior of the buret has been sufficiently rinsed with the titrant solution. Discard the titrant used to rinse the buret and close the stopcock and put your fresh titrant in the buret. Fill the buret as close as you can to the top of the buret. Now open the stopcock fully and allow some of the titrant to be released. This eliminates any bubbles at the tip which can cause error in your data. Try to start your titration with the titrant level near the 0.00 mark. It does not have to be exact so do not waste a lot of your time being exact. Try to get at the 0.00 mark or just below it. Always record this initial volume. The technique for reading the values from a buret are the same as those for the graduated cylinder. The buret is maked by 1 ml lines and the subdivisions represent 0.1 ml. You can therefore, estimate the location of the meniscus to 0.01 ml. 4. Record your initial volume. 5. Place the flask containing the KHP solution on a piece of white paper and position it under the buret. Begin adding your standardized NaOH solution to the flask. Make sure you gently swirl the flask as you are adding the titrant. You will momentarily see a spot of pink color when you add the titrant, but it will disappear as you swirl. When you notice that the pink color is more persistant before being swirled away, you know you are nearing your endpoint. Slow down the addition of titrant at this time. You may want to keep a preliminary final volume reading recorded just in case you overshoot with one more drop. You have reached the endpoint when one drop of titrant causes the whole solution to turn from colorless to a permanent very faint pink color which will not disappear when you swirl. Record this final volume from the buret. If you have a dark pink color, you have gone too far past your endpoint and the data from that titration is invalid. Leave any leftover solution in your buret since you will have to do this another two times. Just refill the buret back to the 0.00 mark or close to it. It is not necessary to have the same initial volume for each titration. Just realize that your endpoint will occur at a slightly different volume if your initial volume is not the same as the first titration. 6. Repeat steps 1-5 with the exception of step 3, two more times. 7. When storing your NaOH solution for next week, make sure you cap your polyethylene bottle very tightly. 8. Calculate the molarity of your NaOH solution for each of the three trials. These values should agree within a range of.005 M. If they do not, repeat the procedure one more time. Calculate the average molarity. 9. Obtain an antacid tablet. Record the mass of this tablet. 10. Grind up the tablet using a mortar and pestle to a fine powder and transfer to a weighing vial. Store the weighing vial in your drawer for next week. DO NOT THROW OUT YOUR STANDARDIZED NaOH SOLUTION!!!!!!!!!!! YOU NEED IT FOR NEXT WEEK!!!!!!!! 6

7 Part II Determining Amount of Calcium Carbonate in a Commercial Antacid Data that you should have written down for EACH titration for Part II are: Titration Reactions Mass of sample of antacid tablet Final Buret Reading Initial Buret Reading Volume of NaOH Molarity of NaOH solution (From Part I) Moles of NaOH added a) Moles of HCl present after adding antacid tablet (determined from moles of NaOH added.) b) Moles of HCl present before adding antacid tablet (25.00 ml of HCl of known molarity) c) Moles of HCl neutralized by the antacid tablet (difference between b) and a)) Moles of CaCO 3 (calculated from moles of HCl neutralized (c)) Mass of CaCO 3 (mg) in sample of antacid tablet Mass of CaCO 3 (mg) in antacid tablet Average mass of CaCO 3 in antacid tablet You should have sample calculations of just ONE titration shown in your laboratory notebook. 1. Weigh out between g of the ground antacid tablet and place in a 250 ml Erlenmeyer flask. 2. Using a ml pipet, add ml of a 0.1M HCl solution into the 250 ml Erlenmeyer flask containing the sample of ground antacid tablet. 3. Swirl the Erlenmeyer flask until the tablet is dissolved. 4. Boil the solution gently for 1-2 minutes to remove the CO 2. You will first see little bubble forming which is the CO 2 bubbling off. When you start to see bigger bubbles form, take the flask off the heat. Avoid boiling off the solution. If it is necessary you can rinse the sides of the flask with distilled water. Wait to do step five until the flask is cool to the touch. 5. Add 5-7 drops of bromocresol green indicator. The solution will turn yellow. Titrate with the standardized NaOH solution from Part I. As you titrate, the solution will turn from yellow to a blue/green (teal) colour and then to blue. When you see the blue/green (teal) colour, it means that you are getting very close to the end point. You are only a few drops away from the blue colour. The blue colour indicates that you have reached the end point. 6. Repeat step one at least two more times. 7. Calculate the moles of HCl originally in the 250 ml flask before adding the antacid tablet. The difference between the moles of HCl originally present and the moles of acid neutralized by the NaOH gives you the moles of HCl neutralized by the antacid. From this number of moles of HCl, you can calculate the number of moles of calcium carbonate that reacted with the HCl and then the amount of mg of calcium carbonate present in the antacid tablet. 7

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