Chem 1312 Handout Experiment ONE Laboratory Time Required Special Equipment and Supplies Objective Safety First Aid GETTING THE END POINT TO APPROXIMATE THE EQUIVALENCE POINT Two hours Balance Potassium hydrogen phthalate, KHP (s) Burets 0.1 M Sodium hydroxide, NaOH (aq) Buret clamp Vinegar ph Electrode Phenolphthalein ph Meter Universal Indicator 10-mL Pipet ph = 7 Buffer Pipet bulb ph = 4 Buffer Stirring bar Magnetic stirrer In performing this experiment, the student will consider the factors that affect an indicator s suitability for use in an acid/base titration. Bases, such as sodium hydroxide, can cause skin burns and are especially hazardous to the eyes. Although vinegar is a dilute solution of a weak acid (acetic acid), it is, nevertheless, advisable to avoid splashing it in the eyes.. Following skin contact with either sodium hydroxide, wash the area thoroughly with water. Should sodium hydroxide (or even vinegar) get in the eyes, rinse them with thoroughly with water. At least 20 minutes of flushing with water is recommended. Then seek medical attention. There are a number of experiments in your lab manual that involve acid/base indicators. In Experiment 8, phenolphthalein is used to mark the end point in the titration of vinegar with sodium hydroxide. Methyl orange serves the same purpose in the titration of hydrochloric acid with ammonia (Experiment 30). In Experiment 13, bromothymol blue changes color when carbon dioxide dissolves in water, creating an acidic solution. In this experiment we explore the question of why different indicators change color at different ph values and illustrate that the end point of a titration (signaled by an indicator s color change) does not necessarily correspond to the equivalence point in the titration. PRINCIPLES The equivalence point in a titration is the point at which the acid and base have been mixed in stoichiometric proportion, meaning that neither acid nor base is in excess. For example, if 50.00 ml of 0.1000 M acetic acid is titrated with 0.1000 M sodium hydroxide, the equivalence point will be reached when 50.00 ml of base have been added to the acid. At that point in the titration, the acid and base have exactly neutralized each other. The question is How can one determine that this point has been reached?
In the early portion of the titration, the chemical amount of base that has been added to the titration mixture is less than the chemical amount of acid present and the ph of the titration is less than 7 (at 25 C). At the equivalence point, the titration mixture is essentially a 0.05000 M solution of sodium acetate and should have a ph of 8.7 (at 25 C), because of the hydrolysis of the acetate ion (see Experiment 22). Further addition of base raises the ph well above 7 (25 C) because there is no longer any acid present to react with the extra sodium hydroxide. These changes in ph take place without any apparent changes in the appearance of the reaction mixture. The equivalence point is revealed either by monitoring the changing ph of the solution (by the use of a ph meter) or by finding an indicator that will change color in the vicinity of the equivalence point. However, not all indicators will change color at the desired. point in the titration. Indicators As discussed previously (see Experiment 21), indicators are weak acids that have different colors at low ph values (when the indicator is predominantly in its HIn form) and at high ph values (when the indicator is predominantly in its In form). A good rule of thumb is that the indicator will display its low ph color when the ratio of [HIn] to [In ] has a value of 10 or more; conversely, the indicator will display its high ph color when the ratio of [HIn] to [In ] has a value of 0.1 or less. Consider an indicator which is blue in the HIn form and is red in the In form. If the indicator has a K a value of 1.0 10 4, it will appear blue at ph s below 3 and red at ph s above 5 (see EquationsONE.1 through ONE.3). K a = [H 3 O + ] [In ] (ONE.1) [HIn] 1.0 10 4 = [H 3 O + ](1); [H 3 O + ] = 1.0 10 3 when [HIn] = 10[In ] (ONE.2) (10) 1.0 10 4 = [H 3 O + ](10); [H 3 O + ] = 1.0 10 5 when [In ] = 10[HIn] (ONE.3) (1) Calculations similar to those shown in Equations ONE.1 through ONE.3 reveal that, if the indicator had a K a value of 1.0 10 8, it would appear blue at ph s below 7 and red at ph s above 9. Thus, an indicator with K a = 1.0 10 4 would change color well before the equivalence point was reached in the titration of acetic acid by NaOH while an indicator with K a = 1.0 10 8 would change color just after the equivalence point had been reached in the same titration. Universal Indicator Universal Indicator is actually a mixture of various indicators, chosen so that the mixture will undergo several color changes as the ph of the solution being titrated varies from a value of 4 to a value of 10. In this experiment, you will perform a ph titration of acetic acid by sodium hydroxide in the presence of Universal Indicator. This will permit you to determine whether a given color change occurs near the equivalence point or not.
PROCEDURE Procedure in a Nutshell Standardize 0.1 M NaOH via titration with KHP, using phenolphthalein as an indicator. Dilute 10.00 ml of vinegar to 100.00 ml. Standardize the resulting 0.1 M HC 2 H 3 O 2 via titration with the NaOH, again using phenolphthalein as an indicator. Add 3 drops of Universal Indicator to 25 ml of the dilute acetic acid and titrate it with NaOH, monitoring the course of the titration with a ph meter. Standardizations Clean a buret and prepare it for use in the standardization of sodium hydroxide according to the directions provided in the Introduction. Accurately weigh, to the nearest 0.1 mg, 0.4 g of KHP. Place the KHP in an Erlenmyer flask, dissolve it in 25 ml of distilled water, and add 2 to 3 drops of phenolphthalein. Read and record the initial volume of NaOH in the buret to the nearest 0.01 ml. Titrate the KHP solution until a faint pink color, which does not disappear when the solution is swirled, is obtained. Read and record the final volume of NaOH in the buret to the nearest 0.01 ml. Refill the buret with sodium hydroxide. Rinse a pipet with two small portions of vinegar. Then, use the pipet to deliver a 10-mL sample of vinegar to a clean (but not necessarily dry) 100-mL volumetric flask. Add 50 ml of distilled water to the vinegar and swirl the flask carefully to mix its contents. Dilute the solution to the 100-mL mark with distilled water; invert the capped flask several times to promote further mixing of the diluted vinegar. Clean a second buret and prepare it for use in the standardization of the dilute acetic acid according to the directions provided in the Introduction. Deliver a 25 ml HC 2 H 3 O 2 sample to an Erlenmyer flask (read and record the initial and final buret readings to the nearest 0.01 ml). Add 2 to 3 drops of phenolphthalein to the acid. Read and record the initial volume of NaOH in the buret to the nearest 0.01 ml. Titrate the HC 2 H 3 O 2 solution to the phenolphthalein end point. Read and record the final volume of NaOH in the buret to the nearest 0.01 ml. Calculate the molarities of the NaOH and the HC 2 H 3 O 2 solutions. Determine the volume of NaOH that will have to be added to 25 ml of HC 2 H 3 O 2 to reach the equivalence point.
ph Titration Prepare the ph meter for titration (set the Function Knob to Standby, the Slope Knob to 100%, and attach the electrode). Immerse the electrode in the ph = 7 buffer solution. Set the ph meter s Temperature Knob to the room temperature (probably 20 C to 24 C). Move the Function Knob to ph and adjust the Calibration Knob until the ph meter reading matches the ph of the buffer exactly. Return the Function Knob to Standby. Rinse the electrode with distilled water and shake it to dry the bulb. Immerse the electrode in the ph = 4 buffer solution. Move the Function Knob to ph and adjust the Slope Knob until the ph meter reading matches the ph of the buffer exactly. Return the Function Knob to Standby. Refill the base buret with sodium hydroxide and the acid buret with acetic acid. Deliver a 25 ml HC 2 H 3 O 2 sample to an Erlenmyer flask (read and record the initial and final buret readings to the nearest 0.01 ml). Add 2 to 3 drops of Universal Indicator to the acid. Read and record the initial volume of NaOH in the buret to the nearest 0.01 ml. Rinse the electrode with distilled water and shake it to dry the bulb. Immerse the electrode in the HC 2 H 3 O 2 solution. Put the stirring bar in the solution. Adjust the magnetic stirrer so that the movement of the stirring bar during the course of the titration does not endanger the ph electrode. Move the ph meter s Function Knob to ph. Record the initial ph of the HC 2 H 3 O 2. Begin the titration. Start by adding 1-2 ml increments of NaOH to the diluted vinegar. Record the ph meter reading after each addition of NaOH. Note the ph of the titration mixture at which each color change occurs. As the volume of base added approaches the calculated equivalence point, decrease the size of the NaOH increment. (Add the NaOH dropwise for 2 ml before and 2 ml after the anticipated equivalence point). Continue adding NaOH and recording the ph (and any color change) until you have gone 10 ml beyond the equivalence point. Prepare a plot of ph (y-axis) versus volume of NaOH added (x-axis). The equivalence point is marked by a large increase in the ph upon the addition of a small volume of base. Calculations Titrations are usually performed to find the molarity of an acid or base. Assume that the molarity of your HC 2 H 3 O 2 solution is correct and calculate the molarity of NaOH (using M av a = M bv b ) using the volume of base added that corresponds to (1) each of the observed color changes and (2) the equivalence point as determined on the ph plot. Calculate the % difference between the molarity based on the ph titration and the molarity based on each color change. State the values of K a required for an indicator to be suitable for use in a titration of HC 2 H 3 O 2 by NaOH. Disposal of Reagents Excess KHP should be placed in the containers used for solid waste. Solutions should be neutralized and diluted. They may then be flushed down the drain.
Date Name Section Desk Number Summary Report on Handout Experiment ONE (take your data in your notebook as usual; Summary Report Sheets are shown only as a guide for the preparation of the data table in your notebook) Standardization of NaOH Solution Mass of KHP and container Mass of container Mass of KHP Final buret reading, NaOH Initial buret reading, NaOH Volume used, NaOH Molarity of NaOH solution Standardization of HC 2 H 3 O 2 Solution Final buret reading, HC 2 H 3 O 2 Initial buret reading, HC 2 H 3 O 2 Volume used, HC 2 H 3 O 2 Final buret reading, NaOH Initial buret reading, NaOH Volume used, NaOH Molarity of HC 2 H 3 O 2 solution
ph Titration Buret reading, ml Volume of base ph color added, ml
Buret reading, ml Volume of base ph color added, ml
Buret reading, ml Volume of base ph color added, ml
Date Name Section Desk Number Pre-Laboratory Exercises for Handout Experiment ONE These exercises are to be completed after you have read the experiment but before you come to lab to perform it. Consider the titration of 40.00 ml of 0.2000 M acetic acid with 0.2000 M sodium hydroxide. Calculate the ph of the titration mixture at the following points in the titration. 1) start (no NaOH added) 2) after the addition of 10.00 ml of NaOH 3) after the addition of 20.00 ml of NaOH 4) after the addition of 30.00 ml of NaOH 5) after the addition of 39.95 ml of NaOH 6) after the addition of 40.00 ml of NaOH 7) after the addition of 40.05 ml of NaOH 8) after the addition of 42.00 ml of NaOH