CM1011 Exp6_Titration_Fa23

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Name ______________________________________ Section __________________ Experiment 6 Acid-Base Titration CM1001 Fall 2023 1 Acid-Base Titration In your text (Chang and Goldsby 7 th Ed): 4.3 Acid-Base Reactions 4.6 Solution Stoichiometry Extra reading: 17.3 A closer look at Acid-Base Titrations 17.4 Acid-Base Indicators Background: Titration is a method of quantitative analysis in which a solution of known concentration (the standard solution or titrant ) is gradually added from a burette to a measured quantity of unknown solution ( analyte ) until the reaction is complete ( equivalence point ). The equivalence point is also called the stoichiometric point, as it is where the number of moles of titrant is in the stoichiometric ratio (from the coefficients in the balanced reaction equation) with the number of moles of the analyte present in the sample. The titrant is typically delivered from a burette, which accurately measures volume. If the analyte is a solid, it must be weighed, then dissolved in the solvent. A small amount of an indicator is added (if applicable), and then the solution of the analyte is titrated. Since the titration will determine the number of moles of analyte, typically the mass of the solid and the number of moles of solid is used to determine the molar mass. If the analyte is a solution, a known volume must be used, a small amount of indicator is added (if applicable), and then the solution is titrated. Since the titration will determine the number of moles of solute in the analyte, typically the volume of the solution and the number of moles of the solute are used to determine the concentration. In acid-base titration, an acid-base neutralization reaction occurs between the titrant, typically a strong base, and the analyte, usually an acid with unknown concentration or molar mass. When the number of moles of the added strong base is at the stoichiometric ratio with the number of moles of the acid, the reaction is complete and the equivalence point is reached. An instrument or a visual change must be used to indicate the endpoint of the titration (when one stops adding titrant). In acid-base titration, an indicator, in very small amount, is added to the analyte solution to indicate the endpoint. An indicator is a compound that have different colors when it is in an acid solution or base solution. When the indicator changes color, the amount of the base added is just slightly more than the amount needed to reach the equivalence point. At this point the reaction is just slightly beyond completion, so the titration should end (thus, the endpoint). If the endpoint is carefully identified, it is very close to the equivalence point. Overview of the Procedure: You will be given a NaOH solution, whose concentration is only known approximately (about 0.1 M). The first step, in Part I, is to standardize the NaOH, i.e. titrate the NaOH solution against a well-characterized acid, to determine the concentration of the NaOH to 4 significant figures. Once this determination is made, your calculated molarity of NaOH is used in the calculations of the titrations in Parts II and III, where the total acidity in a sample of vinegar and a sample of fruit juice are determined, respectively.

Name ______________________________________ Section __________________ Experiment 6 Acid-Base Titration CM1001 Fall 2023 2 Part I. KHP (potassium hydrogen phthalate, KHC $ H % O % , MW 204.2 g/mol) is a weak acid that is used to make solutions of very well-characterized concentrations, and hence it is a useful primary standard . It is a solid under normal lab conditions, it can be highly purified, it does not easily oxidize, and it has a high MW, resulting in good precision and accuracy when weighing convenient sample sizes. KHP is a monoprotic acid, reacting 1:1 mole ratio with NaOH: KHC $ H % O % (aq) + NaOH (aq) à KNaC $ H % O % (aq) + H ) O (l) Sample calculation of a standardization run: A student is given a starting solution of NaOH to standardize, known to be approximately 0.1 M. It took 19.98 mL of the approximately 0.1 M NaOH solution to reach the equivalence point with a solution of KHP that was prepared with 0.4168 g KHP. What is the concentration of the NaOH, to 4 s.f.? 0.4168 g/ (204.2 g/mol) = 2.041 x 10 -3 moles KHP 2.041 x 10 -3 moles KHP x [ 1 mol NaOH/ 1 mol KHP] = 2.041 x 10 -3 moles NaOH 2.041 x 10 -3 moles NaOH/ 0.01998 L = 0.1022 M NaOH Note: You cannot leave out the 1:1 mole ratio in the calculation because that is how moles of KHP are converted to moles of NaOH. Once the concentration of the NaOH titrant is known, the titrant is used to titrate samples of vinegar and fruit juice, to determine the unknown concentration of total acid in each of these samples. Part II. White vinegar will be titrated with the standardized NaOH. White vinegar is a dilute solution of a single acid (acetic acid, HC ) H * O ) , a monoprotic acid), so the data from the titration can be used to determine the concentration of acetic acid in the vinegar. Hydrogen ions from the HC ) H * O ) react with hydroxide ions from the NaOH in a 1:1 mole ratio to produce water and sodium acetate in the overall reaction: HC ) H * O ) (aq) + NaOH (aq) à H ) O (l) + NaC ) H * O ) (aq) A general form for this type of reaction, a monoprotic acid (HA) with NaOH is: HA (aq) + OH - (aq) à H ) O (l) + A , (aq) or HA (aq) + NaOH (aq) à H ) O (l) + NaA

Name ______________________________________ Section __________________ Experiment 6 Acid-Base Titration CM1001 Fall 2023 3 Sample calculation of a run to determine the concentration of a monoprotic acid: A 10.00 mL sample of an unknown monoprotic acid solution is titrated, using the standardized 0.1022 M NaOH solution. It took 20.04 mL of the NaOH to reach the equivalence point. What is the concentration of the monoprotic acid, HA, to 4 s.f.? 0.1022 M NaOH x 0.02004 L = 2.008 x 10 -3 moles NaOH 2.008 x 10 -3 moles NaOH x [ 1 mol HA / 1 mol NaOH] = 2.008 x 10 -3 moles HA 2.008 x 10 -3 moles HA/ 0.01000 L = 0.2008 M HA Note: You cannot leave out the 1:1 mole ratio in the calculation because that is how moles of NaOH are converted to moles of HA. Part III Lemon juice will be titrated with the standardized NaOH. Fruit juices typically have multiple organic acids, such as citric, ascorbic (vitamin C), tartaric, oxalic, malic, etc. However, citric acid ( H * C - H . O / , a triprotic acid) is the dominant acid in lemon and lime juice. The calculations in this titration will assume that the determination of total acidity is due only to citric acid. This approximation will give the maximum amount of citric acid in the lemon juice, or total acidity as equivalent to citric acid. The three hydrogen ions from the citric acid react with hydroxide ions from the NaOH in a 1:3 mole ratio to produce water and sodium citrate in the overall reaction: H * C - H . O / (aq) + 3 NaOH (aq) à 3 H ) O (l) + Na * C - H . O / (aq) A general form for this type of reaction, a triprotic acid ( H * A) with NaOH: H * A (aq) + 3 OH - (aq) à 3 H ) O (l) + A *, (aq) or H * A (aq) + 3 NaOH (aq) à 3 H ) O (l) + Na * A (aq) Sample calculation of a run to determine the concentration of a triprotic acid: A 2.00 mL sample of the unknown citric solution is titrated, using the standardized 0.1022 M NaOH solution. It took 18.67 mL of the NaOH to reach the equivalence point. What is the concentration of the triprotic acid, H 3 A to 3 s.f.? 0.1022 M NaOH x 0.01867 L = 1.908 x 10 -3 moles NaOH 1.908 x 10 -3 moles NaOH x [ 1 mol H * A/ 3 mol NaOH] = 6.360 x 10 -4 moles H * A 6.360 x 10 -4 moles H * A/ 0.00200 L = 0.318 M ࠵? ࠵? A Note: These sample calculations for monoprotic and triprotic acids are for illustration. They are not clues to the actual concentrations in the samples.

Name ______________________________________ Section __________________ Experiment 6 Acid-Base Titration CM1001 Fall 2023 4 Experimental Procedure: Review: Reading a burette, Exp. 1 ( http://www.chem.tamu.edu/class/fyp/mathrev/mr-sigfg.htmL) Fig. 1 Reading a burette Burette Read at the bottom of the meniscus. The smallest division on the burette is 0.1 mL. Therefore, our reading needs to have the estimated digit in the hundredth place. A good volume reading is 20.38 ± 0.02 mL. A reading of 20.35 mL ± 0.05 mL is also acceptable. Look at the pictures of a burette below. Note that the numbers get bigger as you go down the burette. This is different from the beaker or the graduated cylinder. This is because the liquid leaves the burette at the bottom Always read a burette at eye level. The rings on the burette aid in reading the meniscus without parallax error, or incorrect volume reading due to looking above or below eye level, as shown in Fig. 2 below. When reading volume with a meniscus, always read the bottom of the meniscus. Using either a white card with a black bar on it or a plain white card (moving up from underneath) behind the burette will help highlight the bottom of the meniscus, as shown in Fig. 3 below. Fig. 2 Fig. 3 Note that the ring in the back of the buret can be seen, looking up at the 5 mL mark and down at the 7 mL mark.

Name ______________________________________ Section __________________ Experiment 6 Acid-Base Titration CM1001 Fall 2023 5 Fig. 4. Experimental setup ↑ Always wear goggles! – 5 pts safety deduction each time you are asked by the instructor or TA to put your goggles on. NaOH is corrosive and can damage eyes, skin, and clothing. If contact accidentally occurs, rinse the affected parts with large amounts of water. Funnel (optional) for filling. Remove the funnel after the burette is filled. In this experiment, NaOH is the ONLY substance that goes in the burette. Burette clamp. Use to adjust the height of the burette so meniscus is at eye level. Reading at an angle results in parallax error. Erlenmeyer flask, containing a) acid to be titrated b) 2- 3 drops phenolphthalein c) water, as diluent or solvent White paper, to enable better detection of the faint pink endpoint

Name ______________________________________ Section __________________ Experiment 6 Acid-Base Titration CM1001 Fall 2023 6 Preliminaries You should first view the Burette Technique video posted along with this handout on the Brightspace course site. Set up the apparatus as in Fig. 4. Take a large beaker and, with a pencil, label it “WASTE” on the white patch on the side. Use this for on-bench disposal of burette rinses and finished titrations. Take a 150 mL beaker, label it “NaOH”, and use it to get about 100 mL of the approximately 0.1 M NaOH from the stock solution in the hood. Put the waste beaker under the burette instead of an Erlenmeyer flask, and rinse the burette twice, using about 25 mL each time. (Note that the “off” position for the stopcock is when the handles of the stopcock are perpendicular to the long axis of the burette). Swirl the funnel to wash the rinse NaOH down the sides of the burette. Any time you will fill a burette with a funnel, you might have to lift the funnel slightly so the liquid flows freely. Once the burette has been rinsed, with the waste beaker still under the burette (to collect any accidental overflow), fill it to your eye level (must be above the 20-mL mark), lifting the funnel to keep from overflowing. Turn the stopcock full on, to blow the air out of the tip. If there is a residual air bubble, open the stopcock so the liquid flows freely and tap or gently jiggle the burette, to knock the bubbles out. Ask your instructor for help if you are not able to get rid of the bubbles. Once the air bubbles are cleared, fill to back up to your eye level. Your burette is now ready to start titrating. As needed use your “NaOH” beaker to get refills. Part I. Standardization of NaOH solution. 1. The actual molarity of the approximately 0.1 M NaOH must be determined by titration to 3-4 s.f., remember the actual number of s.f. is dictated by the s.f. of the worst (lowest number of s.f.) piece of data. Obtain a small weighing boat (keep this boat to be reused for all the KHP trials). Tare the balance with the boat and weigh approximately 0.4 g of KHP (potassium hydrogen phthalate, KHC $ H % O % , MW 204.2 g/mol) on one of the analytical balances. The sample does not have to be exactly 0.4000 g, but the mass should be recorded to 0.0001 g (ex. 0.3976 g, 0.4010 g, etc.) record the mass on the Data Sheet, Part I, Step 1 . (Do the calculations of the number of moles later, at home). 2. Carefully transfer the KHP sample into a 250 mL Erlenmeyer flask, and use the DI water bottle to rinse any residual crystals into the flask. Add about 50 mL of deionized water and 2-3 drops of phenolphthalein indicator to the KHP solution. Swirl or stir with a glass rod until all of the solid is dissolved. Record the initial volume reading of the burette on the Data Sheet, Part I, Step 2. It is not critical what the starting point is, as long as the meniscus is on the burette scale and above the 20-mL mark, since you will need at least 20 mL for each titration.

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