EXPERIMENT #8 Acid-Base I: Titration Techniques OBJECTIVES: Dispense a precise volume of a solution with a buret Titrate a known volume of acid solution with a standard solution of base Reach a proper end point Calculate the concentration of an unknown acid by titration with a standard solution of base BACKGROUND: An acid is a compound whose formula begins with one or more hydrogens and in aqueous solution produces hydrogen ions, H + (aq). Acids may be weak (partially dissociate in aqueous solution) or strong (completely dissociate in aqueous solution). The strong acids are: hydrochloric acid, HCl(aq); hydrobromic acid, HBr(aq); hydriodic acid, HI(aq); sulfuric acid, H2SO4(aq); nitric acid, HNO3(aq); and perchloric acid, HClO4(aq). All other acids are weak acids. Examples of weak acids include: hydrofluoric acid, HF(aq); nitrous acid, HNO2(aq); acetic acid, HC2H3O2(aq); and potassium hydrogen phthalate, KHC8H5O4. A base is a compound that is capable of reacting with hydrogen ions. Most bases we deal with in aqueous solution contain the OH group and in aqueous solution produce hydroxide ions, OH (aq). When hydrogen ions meet hydroxide ions in aqueous solution, they neutralize each other by combining to form water, H2O(l). This is the basis of an acid-base titration. A titration is a general procedure for accurately measuring the volume of one solution dispensed into and reacting with another. At the equivalence point of an acid-base titration, the number of moles of H + is equal to the number of moles of OH. This stoichiometric relationship forms the basis for quantitative calculations. An acid-base titration is performed to determine the amount of acid or base present in an unknown. If we titrate with standard base, a solution of base of known concentration, we can calculate the amount of acid present; similarly, if we know the concentration of acid, we can calculate the concentration of base. Let's look at the simplest case. We will titrate a strong, monoprotic acid (an acid containing only one acidic hydrogen) with a strong base containing one hydroxide group. Consider the reaction of hydrochloric acid with an aqueous solution of sodium hydroxide. The molecular equation is The total ionic equation is HCl(aq) + NaOH(aq) H2O(l) + NaCl(aq) The net ionic equation is H + (aq) + Cl (aq) + Na + (aq) + OH (aq) H2O(l) + Na + (aq) + Cl (aq) H + (aq) + OH (aq) H2O(l) The molecular equation shows that one mole of HCl reacts with one mole of NaOH. Each mole of HCl supplies one mole of H + (aq); each mole of NaOH supplies one mole of OH (aq). Therefore, for each mole of acid present in an unknown, one mole of base is required; for each mole of base present in an 79 P a g e
EXPERIMENT #8 TITRATION TECHNIQUES unknown, one mole of acid is required. At the equivalence point of an acid-base titration, the moles of hydrogen ion are equal to the moles of hydroxide ion: moles H + = moles OH moles H + = moles HCl moles OH = moles NaOH volume HCl(aq) x molarity HCl = volume NaOH(aq) x molarity NaOH If a known volume of HCl is titrated with a known volume of NaOH of known concentration, the molarity of the HCl can be calculated as follows molarity HCl = volume NaOH(aq) x molarity NaOH volume HCl(aq) We know we have reached the equivalence point when an indicator, some substance we have added to the solution being titrated, changes color. In the present experiment the equivalence point will be signaled by a color change from colorless to pink using phenolphthalein as the indicator. Example: 22.37 ml of hydrochloric acid of unknown concentration was dispensed by buret into a flask. A few drops of phenolphthalein were added and the colorless solution was titrated with standard sodium hydroxide solution, 0.1755 M. The pink equivalence point was reached when 31.23 ml of sodium hydroxide solution was added. Calculate the concentration of the HCl solution. molarity HCl = volume NaOH(aq) x molarity NaOH volume HCl(aq) molarity HCl = 0.03123L x 0.1755M 0.02237L = 0.2450 M 80 P a g e
EXPERIMENT #8 TITRATION TECHNIQUES PROCEDURE: (Work independently.) Sodium hydroxide and hydrochloric acid are classified as corrosives and can easily do irreversible damage to your eyes. Please make sure you wear your safety goggles throughout the experiment. 1. Solutions The acid and base solutions have been prepared; the base solution has been standardized. The concentration of the sodium hydroxide solution is listed on the container; be sure to record this value on your data sheet. Take only minimum amounts of acid and base to do the experiment. If you need more, you can always get a refill. Never put material back into a standard solution container. Use clean and dry 150 ml beakers to collect the acid and base. (Your instructor will show you the technique to collect standard solutions if the beakers are not dry.) Carefully label each beaker. Fill the burets by adding the contents of the beaker directly to the buret. 2. Burets a. Find a clean, functional buret. Make sure the stopcock is attached properly, does not leak, and allows a rapid stream to flow when fully opened. Rinse the buret with deionized water. If droplets of water remain inside the buret, the buret must be thoroughly cleaned with detergent solution and a buret brush followed by several thorough rinses with deionized water. b. Add 3-5 ml of the base solution to the buret, turn the buret to an almost horizontal position with the open end over an empty beaker or the sink, and rotate the buret so that the inside surface of the buret is covered with solution. Then drain the solution through the stopcock into the beaker or sink. Repeat the process two more times. c. Fill the buret with the standard base solution. Remove any air bubbles by allowing the filled buret to drain into a beaker with the stopcock fully opened. Difficult-to-remove air bubbles are eliminated by applying a sharp vertical, downward motion to the buret with your wrist. Your instructor will demonstrate. Once the stopcock and buret tip are free of air bubbles, place the buret securely in a buret clamp, fill the buret with the acid solution slightly above the zero mark, and adjust the meniscus close to, but below, the zero mark using a white background to clearly define the bottom of the meniscus. 3. Titration a. Using a volumetric pipet transfer 25 ml of the HCl solution into a clean 125 ml Erlenmeyer flask. b. Add two drops of phenolphthalein indicator to the flask. Place a white background under the flask in order to see the end point clearly. c. Slowly add titrant to the flask, while gently swirling its contents. As the sodium hydroxide solution is added, a pink color appears where the drops of the base come in contact with the solution. The color disappears with swirling, but near the end point the color disappears more slowly. Near the 81 P a g e
EXPERIMENT #8 TITRATION TECHNIQUES end point the titrant must be added drop by drop or even one half drop by one half drop. It is most important that the flask be swirled constantly throughout the entire titration to ensure homogeneity. The faintest pink color that persists in front of a white background for 30-60 seconds after swirling indicates the end point has been reached. Reread the buret after this time frame and record on your data sheet. Any drop adhering to the buret tip should be added to the flask. Pay close attention as your instructor demonstrates the technique. Avoid splashing during the titration. Titrant solution on the sides of the flask can be rinsed with deionized water from a wash bottle. Your instructor will demonstrate the technique of adding a partial drop to the flask. d. Record the initial and final reading on the buret containing the standard base solution. e. Compute the molarity of the standard acid solution. f. Repeat steps [3a to 3e] until a consistent value (good precision) for the molarity of the HCl solution is calculated. What are the buret readings when the volumes in the burets are as shown below? FIGURE I: Correct use of a buret 82 P a g e
NAME Section Date DATA AND CALCULATIONS: Titration Techniques Concentration of the Standard Base Trial Volume of Acid, ml Buret Reading Volume of Base, ml Calculated Acid Molarity I Final Volume Used II Final Volume Used III Final Volume Used Average Molarity CLEARLY show how you calculate the acid concentration for each trial. Set up the equations, substitute the numbers, rearrange to solve for the unknown, and then calculate the unknown value. 83 P a g e
Trial Volume of Acid, ml Buret Reading Volume of Base, ml Calculated Acid Molarity I Final Volume Used II Final Volume Used III Final Volume Used Average Molarity CLEARLY show how you calculate the acid concentration for each trial. Set up the equations, substitute the numbers, rearrange to solve for the unknown, and then calculate the unknown value. 84 P a g e
NAME Section Date ADDITIONAL ASSIGNMENT I: Titration Techniques 1. How many moles of sodium hydroxide are there in each of the following solutions: Clearly show each calculation and make proper use of significant figures. (a) 1.250 L of a 0.2087 M solution? (b) 0.5330 L of a 0.2087 M solution? (c) 233 ml of a 0.2087 M solution? 2. Calculate the concentration (in molarity) of a NaOH solution if 25.00 ml of the solution are needed to neutralize 17.42 ml of a 0.3125 M H2SO4 solution. First write a properly balanced chemical equation. 85 P a g e
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NAME Section Date ADDITIONAL ASSIGNMENT II: Titration Techniques 1. 7.02 ml of 6.35 M NaOH are diluted with water to a volume of 400. ml. We are asked to find the molarity of the resulting solution. a. First calculate the number of moles of NaOH in 7.02 ml of 6.35M NaOH. b. Since the total number of moles of NaOH is not changed on dilution, the molarity after dilution can be found by using the final volume of the solution. Calculate that molarity. 2. In an acid-base titration it takes 26.52 ml of 0.1052 M NaOH solution to neutralize (reach the end point) of 27.46 ml of HCl. a. write a balance chemical equation for this reaction b. find the moles of NaOH needed to reach the equivalence point c. find the molarity of the HCl solution. 87 P a g e
3. If 8.00 ml of 6.0 M NaOH is diluted with distilled water to a volume of 500.00 ml, what is the molarity of the resulting solution? 4. What volume of a 0.500 M HCl solution is needed to neutralize each of the following: You need to write a balanced chemical equation to help you visualize the stoichiometry. (a) 10.0 ml of 0.300 M NaOH solution (b) 10.0 ml of a 0.200 M Ba(OH)2 solution 88 P a g e