Volumetric Analysis Lab for General Chemistry Lab

Volumetric Analysis¾Acid-Base Titrations Introduction +VOLUMETRIC ANALYSIS involves the application of the MOLE CONCEPT to reactions of solutes in SOLUTION. In most lab practice, CONCENTRATION of a solution is most conveniently expressed in terms of the solutionÕs MOLARITY, M: the number of moles of solute dissolved per liter of solution. From the MOLARITY, the amount of solute involved in a reaction in solution can be determined by just measuring the VOLUME of the solution in question. For example, 50.0 mL of 0.l00 M salt solution contains (mole = M x L): (0.0500 liters)(0.l00 moles/liter) = 0.00500 moles of salt +Volumetric analysis employs the technique of TITRATION in which the VOLUMES of solutions required to react EXACTLY with each other are determined. The VOLUMES of the reacting solutions are measured carefully with a device called a BURET. The scale divisions on the buret enable you to measure volumes to the nearest 0.01 mL. The point where one solution has been added in just the right amount to react exactly with the other solution present is called the EQUIVALENCE POINT of the titration. The equivalence point is determined visually by use of a substance called an INDICATOR (in acid-base titrations, such as in this experiment, the indicator used is a dye that CHANGES COLOR if the solution turns from slightly acidic to slightly basic). +In this experiment, you will be given an UNKNOWN SOLUTION containing a certain concentration of ACETIC ACID, CH3COOH. We will determine the CONCENTRATION of acetic acid present, by titrating the UNKNOWN acetic acid solution with a solution of SODIUM HYDROXIDE (NaOH) of KNOWN concentration. Acetic acid and sodium hydroxide react according to CH3COOH + NaOH ˆ CH3COONa + H2O i.e., in a one-to-one ratio. Suppose we find it takes 25.02 mL of 0.1000 M NaOH solution to titrate 22.09 mL of an unknown acetic acid solution. The number of moles of NaOH that have been used in the titration is moles NaOH = 0.02502 liters x 0.l000 moles/liter = 0.002502 moles NaOH But, since the chemical reaction has 1:1 stoichiometry, there must have been 0.002502 moles of acetic acid in the unknown solution sample. Therefore, the CONCENTRATION in moles/liter, of the acetic acid in the unknown must be There is one minor complication to this experiment that we have to get around, however. SODIUM HYDROXIDE SOLUTIONS ARE NOT VERY STABLE (they react with the glass of the bottles used to store them, and also can absorb carbon dioxide gas from the air). So we canÕt provide you with a solution of NaOH whose concentration is known precisely enough for the accuracy required in this sort of experiment. You will have to prepare your own solution of NaOH, and then DETERMINE its concentration accurately and precisely. Sodium hydroxide also reacts on a 1:1 stoichiometric basis with HYDROCHLORIC ACID, HCl NaOH + HCl ˆ NaCl + H2O + (net ionic: H + OH - ˆ H2O) Luckily, solutions of hydrochloric acid can be made up conveniently to KNOWN CONCENTRATIONS, and can be KEPT STORED for long periods of time without significant decomposition. In this experiment, you will be provided with a STANDARD solution of exactly 0.1000 M HCl solution. You will prepare a solution of sodium hydroxide of APPROXIMATELY 0.1 M, and then, by titrating the NaOH solution with the STANDARD HCl solution, you will determine the EXACT concentration of the NaOH solution (to four significant figures). The calculations involved are completely analogous to the calculation described above for the unknown acetic acid. Once you know the concentration of your own NaOH solution accurately, you can then use this solution to titrate the acetic acid unknown in order to determine its concentration. +In addition to determining the concentration of the acetic acid unknown solution in terms of its MOLARITY, we also want to report the amount of acetic acid present in the unknown solution as the PERCENT acetic acid in the solution BY WEIGHT. To convert from the MOLAR CONCENTRATION (of a solution) to a PERCENTAGE (of solute by weight), the DENSITY of the solution is needed for your unknown (which is mostly water). You can take the density of the solution to be 1.00 g/mL. Additionally the FORMULA WEIGHT of the solute (CH3COOH) is equal to 60.0 g/mole. +Suppose we do a titration as described above, and determine that our acetic acid unknown has a concentration of 0.1133 moles/liter. The following would be how we could convert this MOLARITY to PERCENT BY WEIGHT. Notice how the units cancel: (Notice that the units are Òg acetic acid/g solution.Ó) The factor of 100 converts the FRACTION to PERCENT (per 100). Numerically, the conversion from molarity to percent by weight is trivial, since many of the numbers in the above conversion cancel out. Summary Approximately 0.1 M NaOH solution is prepared by dilution of concentrated stock solution (2.0 M). The approximately 0.1 M NaOH is standardized by titration with standard 0.1000 M HCl solution using phenolphthalein as an indicator (this indicator is colorless in acidic solution, but pink in basic solution). The now known NaOH solution is then used to determine the concentration of an unknown acetic acid solution by a similar titration. The percent by weight of acetic acid in the unknown is calculated. Supplies Two burets and clamp; stock NaOH solution; standard 0.l000 M HCl solution; phenolphthalein indicator solution; acetic acid unknown. RECORD THE NUMBER OF THE UNKNOWN IMMEDIATELY IN YOUR LAB NOTEBOOK. CHEMICALS Hydrochloric acid (HCl)¾toxic and corrosive, will burn skin Sodium hydroxide (NaOH)¾toxic and corrosive, easily absorbed through the skin Acetic acid¾corrosive to skin and tissue, can cause skin burns Phenolphthalein¾no major health risks Procedure NO SOLUTIONS FROM THIS EXPERIMENT SHOULD GO DOWN THE DRAIN. CAUTION! WEAR SAFETY GLASSES AT ALL TIMES!! (1) Preparing the burets For accurate results in any titration, the burets used must be scrupulously CLEAN. Water (or titrating solutions) should flow IN SHEETS down the inside of the buret, and should NOT ADHERE in drops to the inside of the glass. Obtain two burets and a buret clamp from your labroom. Rinse the burets thoroughly with water to remove any previous chemicals. If water adheres to the interior of the buret, obtain some soap solution and a special Òburet brushÓ from the stockroom, and scrub the buret until water no longer beads up. Rinse finally with water. Before a buret is subsequently filled with any solution, THE BURET SHOULD BE RINSED WITH THREE OR FOUR 5 to10 mL PORTIONS OF THE SOLUTION THAT ULTIMATELY IS TO BE PLACED IN THE BURET. This makes sure that the solution to be placed in the buret will not be diluted inadvertently by any residual water left in the buret from the washing. (2) Preparation of solutions and filling the burets Obtain 25 mL of 2.0 M NaOH solution, and add this to 475 mL of distilled water in your Florence (or 500 mL Erlenmeyer) flask. Stir the solution thoroughly. This will produce 500 mL of approximately 0.l0 M NaOH solution. Stopper the flask when not using the solution (to keep the NaOH from reaction with carbon dioxide in the air). Thoroughly CLEAN and DRY a 250 mL beaker, and obtain around 75 mL of 0.1000 M HCl standard solution. Keep the beaker covered with a watch glass during the experiment (the beaker must be clean and dry, and must be kept covered, to prevent the concentration of the HCl solution from changing). Rinse one buret several times with small portions of your dilute NaOH solution. Discard the rinsing into a 400 mL beaker from your locker to be used for waste. Finally fill the buret with the dilute NaOH solution. Rinse the other buret several times with small portions of the standard HCl solution Discard the rinsing into a 400 mL beaker from your locker to be used for waste. Finally fill your buret. Make certain there are no air bubbles trapped in the tip of either buret. (Ask your lab instructor if you canÕt get the bubbles out yourself.) Record (in your lab notebook and on the report sheet) the initial readings of both burets, reading across the bottom of the curved liquid surface (called the MENISCUS). Make your readings to TWO DECIMAL PLACES (any less precision will give meaningless results). NOTE: +Initial readings should be reported as zero mL (or close to it), not 50 mL. (3) Standardization of NaOH solution Run around 20 to 25 mL of the HCl standard solution from the buret into a clean 250 mL Erlenmeyer flask, and add 2 or 3 drops of the phenolphthalein indicator solution. It is NOT necessary to read the burets at this point. The 20 to 25 mL of solution run out at the start will be INCLUDED in the reading taken at the end point (color change) of the titration. Titrate the HCl solution by adding NaOH solution from the other buret until the solution in the Erlenmeyer flask just turns pink (keep swirling the flask while adding the NaOH). If the pink color does not remain for at least 30 seconds, add another drop of NaOH until it does stay pink for at least 30 seconds. The appearance of the pink color means you have added MORE NaOH than is required to react with the HCl in the flask. Carefully Òback titrateÓ the solution in the flask by adding HCl from the other buret drop by drop (swirling the flask constantly), until ONE SINGLE DROP of HCl from the buret causes the pink color to JUST disappear. Confirm that you are now at the end point by adding ONE SINGLE drop of NaOH from the buret (careful!); the pink color should just return faintly. Record the readings of the burets (again reading the bottom of the meniscus) to the nearest 0.01 mL. (Final readings should be read somewhere between 20 and 25 mL.) Discard this solution into your waste beaker. When the waste beaker becomes filled, discard the contents into the WASTE ACID-BASE container. From the volume of each solution used, and from the known molarity of the HCl, calculate the molarity of the NaOH solution. Refill the burets and repeat the entire process above. Calculate the molarity of the NaOH from the results of this second titration. If the two molarities of NaOH do not agree within 1%, repeat the standardization a third time or until two molarities are obtained within 1% of each other. Use the AVERAGE of your two closest molarities as the concentration of your NaOH solution. Note: It is not necessary to always fill a buret to the zero mark (especially if your height prevents you from seeing the zero mark easily). If the actual initial volume reading is recorded, as well as the final volume reading (at the color change), then the volume of solution used is just (Volfinal Volinitial). (4) Analysis of the acetic acid unknown Drain the buret that contained HCl into your waste beaker, rinsing it several times with distilled water. Discard the rinses also into your waste beaker. Then rinse the buret with several small portions of your acetic acid unknown. Again, discard the rinses into your waste beaker. Finally fill the buret with the unknown, again making sure to avoid air bubbles. Run about 20 to 25 mL of the unknown from the buret into a clean 250 mL Erlenmeyer flask, add 2 or 3 drops of phenolphthalein solution, and titrate with the standard NaOH to the end point. Discard this solution into your waste beaker. Based on the volumes of NaOH and unknown used, and on the now known molarity of the NaOH, calculate the molarity of the unknown acetic acid solution. Repeat the titration of the unknown until you get two calculated molarities of the unknown within 1% of each other. Use the average of your two best values as the concentration of your unknown. Discard the titration solutions, any leftover solutions in your burets into your waste beaker. Transfer the contents of your waste beaker and any leftover NaOH (in your Florence flask or 500 mL Erlenmeyer flask) into the WASTE ACID-BASE container. Finally, calculate the percent by weight of acetic acid in your unknown. Since your unknown solution is quite dilute, you are valid in making the approximation that its density is 1.00 g/mL; the molecular formula of acetic acid is CH3COOH, which means a formula weight of 60.0g/mole. REMEMBER TO TURN IN THE YELLOW COPY OF YOUR LAB NOTEBOOK TO YOUR TA. Acid-Base Titrations Report Sheet (30 Points) Name ______________________________ +Lab Day ____________ Lab Instructor_____________________ +Date _______________ Show calculations on the back of the page. Molarity of standard HCl solution _________________ M Unknown no.__________ 3. +Standardization of NaOH solution Buret Readings Final Volume, mL Initial Volume, mL Volume Used, mL Molarity NaOH Average Molarity + First + Trial +Second +Trial +HCl +NaOH +HCl +NaOH 4. +Titration of Acetic Acid Unknown Buret Readings + First + Trial +Second +Trial +CH3COOH +NaOH +CH3COO +NaOH H Final Volume, mL Initial Volume, mL Volume Used, mL Molarity CH3COOH Average Molarity Percent acetic acid of your unknown. (Show work on back of sheet or separate piece of paper.) +% acetic acid _________________ If additional titrations are needed, use the table below for your data. Buret Readings + Third + Trial +Third +Trial +HCl +NaOH +CH3COO +NaOH H Final Volume, mL Initial Volume, mL Volume Used, mL Molarity CH3COOH Average Molarity Acid-Base Titrations Post-Lab Questions (40 Points) Name ______________________________ Lab Instructor_____________________ 1. +Define equivalence point. +Lab Day ____________ +Date _______________ 2. +What is an indicator? 3. +Why does phenolphthalein turn pink in the titrations you performed? 4. +Why is it necessary to experimentally determine the molarity of the NaOH solution used in +the titrations in this experiment instead of just using the value calculated by your dilution? 5. +Write a balanced net ionic equation for the reaction that occurs between acetic acid and +sodium hydroxide.