Learn to do quantitative titration reactions. Observe the mole ratios of several simple chemical reactions.

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1 CHAPTER 6 Stoichiometry of Reactions in Solution Objectives The objectives of this laboratory are to: Learn to do quantitative titration reactions. Observe the mole ratios of several simple chemical reactions. Background This semester you have learned about balancing chemical reactions, and making calculations from those balanced reactions. One of the important ways that these skills are used in chemistry and in many other sciences is in quantitative analysis. The discipline of quantitative analysis always begins with a balanced equation for a process, and then calculates quantities of standards and unknowns using that balanced equation. The measurements and calculations in quantitative analysis are usually based on the mass of a known substance; this substance is called a standard, and is always very carefully chosen to minimize errors and inaccuracies. You will not have to prepare the standards for this experiment; it has been done for you by the stockroom staff. In order to make calculations, one must have some means of telling when a reaction is complete. In this experiment we will introduce you to one method of determining when a reaction is complete. We will use an indicator, which changes color at the end of a reaction. You will investigate balanced reactions between four different substances, some of which may be familiar. Hydrochloric acid,, is a common strong acid, often used in household cleaners such as toilet bowl cleaners. Sulfuric acid,, is not often used in the home, but it is one of the top ten industrial chemicals in quantity produced and used. Another substance you will use is sodium hydroxide,, whish is commonly referred to as lye. It is a common active ingredient in household drain cleaners. The last of the four compounds you will use is barium hydroxide,. This chemical is not often used in the household or in industry. The first two of these substances are acids, and the second two are bases. You will spend time later this semester discussing the chemistry of these two important classes of compounds. In this experiment, we will not be investigating the acid/base chemistry 137

2 Chemistry 121 Chemistry involved, but will use these reactions to investigate the stoichiometry of reactions. The four possible reactions between these components are: + O + NaCl + O + Na 2 + O + Ba + O + BaCl 2 As part of your pre-lab preparation, you should balance each of these unbalanced equations. Notice that while water ( O) is a product in all of these reactions, it is not a reactant. The water in which the reactants are dissolved does not participate in the reactions, although a little additional water is generated by the reactions as they occur. Acid and base titrations are some of the most common procedures chemists perform in their every day work. Each of the above reactions will be performed in turn, measuring the quantities of each reactant and watching for the indication of when the reaction is complete. You will then perform calculations to discover how many moles of each reactant were used in order to experimentally determine the stoichiometry of the reaction. 138 At this point in the semester, you may or may not have explored the concept of molarity. In simple terms, molarity is merely a unit of concentration. In a one molar solution (denoted 1 M), there is precisely one mole of solute dissolved in exactly 1 L of total solution. For example, a 0.1 M solution of (a commonly used concentration) will contain 0.1 moles of dissolved in 1 L of solution. You will use the molarities of these solutions to calculate the number of moles of reactants. If you know the molarity (moles/l) and the volume used (converted to L!), you can use simple unit analysis to calculate the number of moles present. Safety Precautions Wear your safety glasses at all times during the lab, and be careful not to rub your eyes with your hands when you have traces of any of the solutions on them. The solutions are mild skin irritants and can damage clothing. Rinsing your hands or clothes off with tap water, warm or cold, will remove all traces in just a few seconds. Materials Equipment Buret (check out) 10 ml volumetric pipette Buret clamp, beakers, and wash bottle (drawer) Beakers or Erlenmeyer Flasks (drawer) Reagents: (write down the standardized concentrations from the boxes) phenolphthalein indicator solution

3 Stoichiometry of Reactions in Solution Chapter 6 Procedures Titrations A titration is an experiment in which a measured volume of one reactant is placed in a reaction vessel (in this case, an Erlenmeyer flask), and the second reactant is added slowly to the first from a buret. The volume of the second reactant can be read at any time from the buret so the quantities of the reactants are always either known or calculable. An indicator of some kind is used to mark the point at which the quantities of the two reactants are just equivalent; this is the end point. In other words, at the end point, you have added exactly enough titrant (the reactant in the buret) to react completely with the reactant in the flask. You are using the indicator phenolphthalein, which changes color when the reaction is complete. Carrying Out the Titration You will use the buret that your TA has described to you at the beginning of the class period. The apparatus must be thoroughly rinsed three times with small amounts (about 5 ml) of the titrant solution. For this first titration that will be the solution. Your TA will demonstrate the technique for rinsing the buret safely and properly. Be sure to measure your solutions out into smaller containers and do not try to fill the buret directly from the stock bottle! Instead, get about 50 ml of solution in a clean glass beaker and bring it back to your work area. You will then fill the buret from the smaller beaker using your funnel. It is very important during this lab to be sure to label all beakers at your station. After rinsing your buret with the solution, clamp it securely to the ring stand. Fill the buret with the solution to slightly above the zero mark. Put a new beaker under the apparatus, open the stopcock, and run a small amount of solution out of the buret to remove any air bubbles from the stopcock and tip. Keep the beaker around so that it may serve as a waste receptacle, marking waste on the outside of it. Using your 10 ml volumetric pipette (your TA will demonstrate the proper use of a volumetric pipette), measure out precisely ml of into a 50 ml Erlenmeyer flask and add 2 drops of phenolphthalein indicator. Add approximately 15 ml of deionized water to the. Position the beaker containing the solution under the buret. Record the initial volume on the buret to the correct number of significant figures (you do not have to start at 0.00 ml!). Open the stopcock, adding solution to the solution, gently swirling the beaker to mix the solutions. The titration will require around 7 ml of titrant, so you can add the first 5 ml of titrant quickly. Slow the addition of the. You will notice the solution turns pink but this color fades on swirling as you approach the end point of the titration. Adding the dropwise with swirling, titrate to a faint pink end point (the end point is sharp the solution will remain pink on the addition of one drop of the at the end point). Record the final volume on the buret. Repeat the titration. 139

4 Chemistry 121 Chemistry For the second titration, perform it exactly in the same manner, but use instead of. Be sure to thoroughly rinse your volumetric pipette to remove any trace that remained in the pipette. The end point of this titration will occur at approximately 16 ml. Do this titration twice. For the last two titrations, you will replace the in the buret with. Be sure to rinse the buret with barium hydroxide to avoid contamination. Measure precisely 10 ml of and add deionized water and phenolphthalein into your Erlenmeyer flask as before. Titrate with as you did with. This titration will take approximately 5 6 ml. Repeat this titration. Perform the last titration in the same manner as the previous one, using instead of. This titration will take approximately ml. Perform this titration twice. Calculations For each of the four titrations, you should have a concentration of acid and of base recorded on your data sheet, and a volume of each reactant. The volume of each acid ( and ) is how much you originally measured into the beaker with the volumetric pipette. The average volume of base ( and is the average of the two titrations. Spaces are provided on the data sheet to record these quantities for each titration. For each of the titrations, calculate from the raw data the moles of acid and the moles of base reacted. All these calculations are similar, so you only need to show how you are doing the calculation once. Record the results of the calculations on the data sheet, including the ratio of (moles acid)/(moles base) for each reaction. An example can be found below. (Note that your measurements are in ml, however, they must all be converted to L to use in these calculations.) (Concentration of ) (volume of ) = ( mol/l) ( L) = moles (avg. concentration of ) (vol. ) = (.101 mol/l) ( L) = moles The volume of is based upon the fact that you measured precisely 10 ml into your flask to begin the reaction. The volume of is derived from your buret readings. ratio acid to base = / = 1.01 ratio acid to base = 1 / 1 = 1 The theoretical ratio is found from the balanced chemical reaction: O NaCl The coefficients in front of and are both

5 Stoichiometry of Reactions in Solution Chapter 6 CHAPTER 6 Name ID No. Worksheet Instructor Partner s Name (if applicable Course/Section ate (of Lab Meeting) Data Summary Concentrations of Solutions: : M : M : M : M Titration of + of : Titration of + of : Titration of + of : 141

6 Chemistry 121 Chemistry Titration of + of : Show calculations for one titration experiment here: 142

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