Answer Happy

INTRODUCTION Some solutions, called buffers, are very resistant to the pH changes normally caused by the addition of an acid or base. These solutions always contain both a weak acid and its conjugate base (or a weak base and its conjugate acid). The solution prepared to find Ka of a weak acid is an example of a buffer solution that happened to contain equal moles of the weak acid HA and the anion A. If a small amount of strong acid were added to that solution, the H+ ion would tend to react with the Aion present, keeping the [H +] about where it was before the addition. Similarly, a small amount of strong base added to that solution would react with the HA present, producing some Aion and water, but not appreciably changing the [OH-]. If similar amounts of acid or base were added to water, the pH would be changed by several units. A more usual way to prepare a buffer is to dissolve the salt of a weak acid (for example, sodium acetate) in a solution of the weak acid (acetic acid in this case).
The pH of the buffer will be equal to the pka (- log Ka) if the moles of HA and Aare equal. By adjusting the molar ratio of HA to A-, the pH can be varied as shown below: Ka [H [A] [HA] = [H*] = K [HA] [A] Taking the log of both sides of the equation and multiplying both sides by -1 gives: [HA] - log H = -log K. - log [A] or: [A] pH = pK, + log [HA] By adjusting the [A]/[HA] ratio, the pH can be made to be either smaller or larger than the pka. It is usually not practical to adjust this ratio by more than a factor of 10 which makes the adjustment +1 pH unit. If the buffer pH needs to be more than 1 unit different from the pka, it is better to find another acid with a more compatible pka value.
In this exercise you will prepare a buffer, measure its pH and test to see how this solution is resistant to pH changes. In the second part of this experiment, the equivalent mass (weight) of an acid will be determined. A weighed sample of an acid is titrated with a standard NaOH solution. The equivalent weight of an acid is the mass which furnishes 1 mole H+ ions, therefore: EW = grams acid moles H The equivalent weight of an acid may or may not equal the molar mass (MM) of the acid. The reason is almost, but not quite, obvious. Let us consider three acids, with the molecular formulas HX, H2 Y and H3Z: one mole HX —one mole H+ therefore, MM = EM one mole H2Y → two moles H+ therefore, MM = 2 x EM one mole H3Z — three moles therefore, MM = 3 X EM H+
PROCEDURE PARTI: BUFFERS Preparation and pH of a 1-to-1 Buffer: Weigh out 1.0 - 1.1 g of Anhydrous Sodium Acetate on a top-loading balance using a tared 250-ml beaker. Calculate in your notebook the exact number of moles of Acetate ions the beaker. Next, obtain the exact molarity of the Acetic Acid you will be using to make the buffer, and then calculate what volume of this - 0.1 M HC2H3O2 is needed so that the moles of Acetic Acid EQUAL the moles of Acetate. Measure this volume of Acetic acid using a graduated cylinder. Dissolve the salt in the acid in the 250- ml beaker. This is now a solution in which moles Ac is equal to moles HC2H302 and: [A] pH = pK + log [HA] is reduced to: pka pH = Look up the Ka value for acetic acid.
Measure the pH of the buffer. Does the measured pH equal the pKą? If not, provide an explanation. Preparation and pH of a 3-to-1 Buffer: Weigh out another sample of NaC2H302 of twice the amount used the first time and dissolve in the 1-to-1 buffer in the 250-ml beaker. Since two more amounts of Acetate are added to the 1- to-1 buffer, the ratio of moles Acetate to moles Acetic Acid is now 3-to-1. The Henderson- Hasselbalch now would be: pH = pka + Log (3.0) Calculate this pH for the combined amount of NaC2H302 Test this buffer solution to see if the measured pH is equal to the calculated value.
Comparison of buffer with water: Obtain two 150-ml beakers and label one as "A" and one as "B". To each add 50.0 mL of distilled water. Add 3 drops of 6 M HCl to the distilled water in Beaker "A". Measure the pH. Add 3 drops of 6 M NaOH to the distilled water in Beaker "B" and measure the pH. Discard these solutions, and rinse and dry the beakers. Using these clean, dry beakers, add 50.0 mL of your 3-to-1 buffer to each beaker. Add 3 drops of 6 M HCl to the buffer solution in Beaker "A" and measure the pH. Add 3 drops of 6 M NaOH to the buffer solution in Beaker "B" and measure the pH. Continue using these solutions for investigation of the buffer's capacity below. Buffer Capacity: Test the range of effectiveness against an added acid by adding more 6 M HCl in 0.25 mL (5 drops) increments to Beaker "A" containing the HCl and buffer from above. Measure the pH after each 5-drop addition. Continue to add HCI until the buffer pH changes by at least 2 pH units from the original pH. Record the total volume of HCI used including the original 3 drops.
Test the range of effectiveness against added base by adding more 6 M NaOH in 5-drop increments to the buffer in Beaker "B". Keep adding the NaOH until the pH changes by at least 2 pH units from the original pH of the 3-to-1 buffer. Record the total volume of NaOH used.
DATA SECTION PARTI: • Top balance #3 • 1.03 g anhydrous sodium acetate . pH meter #504 . 0.1014 M acetic acid • 6 M NaOH • 6 M HCI # pka for acetic acid 4.74 (24-12-24. The Merck Index Online; Royal Society of Chemistry, 2013; M1111) # Since we do not have regular access to the CRC or Merck Index as we typically would in our laboratory setting, the pka value was looked up and provided here with corresponding reference. As a reminder, always include the proper reference whenever making citations like this in your laboratory notebook/report.
Table 1: Buffer Capacity of DI Water Initial pH pH after 3 drops of 6 M Solutions 2.35 (HCI) Beaker "A" 5.39 Beaker "B" 5.74 11.63 (NaOH) Table 2: Buffer Capacity of 1:1 and 3:1 Solution 1:1 Solution 3:1 Solution Mass of sodium acetate 1.03 g 2.06 g pH 4.41 4.89
Table 3: PH Buffer Data for 3:1 Buffers with HCI and NaOH Number of Drops Added 3 pH for "A" with pH for "B" with HCI NaOH 4.80 4.91 5 4.67 5.07 5 4.53 5.57 07 5 4.39 6.32 5 4.26 11.35 5 4.12 5 3.96 07 5 3.55 5 3.25 5 2.51
PART II: Equivalent Mass Weigh the vial containing the unknown solid aid on the analytical balance. Carefully pour out about one third of the sample into a clean, but not necessarily dry, 125 mL Erlenmeyer flask. Weigh the vial with the remaining acid. Add about 25 mL of distilled water and two or three drops of phenolphthalein to the flask. The acid may be relatively insoluble, so don't worry if it all doesn't dissolve immediately. It will dissolve during the titration. Fill one buret with the standardized NaOH solution. Add the standard HCl to another buret. Read both levels carefully and record them. Titrate the solution of the solid acid with the standard NaOH. If the acid appears to be insoluble, add NaOH until the pink color persists, and then swirl to dissolve the solid. If you pass the endpoint, titrate with HCl to get the solution back to colorless. The final pink endpoint should appear on addition of one drop of NaOH. Record the final levels in the NaOH and HCl burets to +/- 0.01 mL.
Pour about half of the remaining solid from your vial into a clean 125 mL Erlenmeyer flask, and re- weigh the vial. Titrate this sample of acid as before with the NaOH and HCl solutions. Titrate the remaining acid. If you use HCl in these titrations, and you probably will, the calculations needed are a bit complicated. To find the number of moles of H+ ion in the solid acid, you must subtract the number of moles of HCl used from the number of moles of NaOH. For a back-titration (which is what we have in this case): moles H+ in solid acid = (moles OH- from NaOH solution) – (moles H+ from HCl solution)
PART II: • Analytical balance #1 • Unknown #3 (white solid granuales with mostly fine grains and some clumping clusters; Maleic acid) • Base buret #20 • Acid buret #18 . 0.1025 M NaOH • 0.1096 M HCI • All three end points very faint pink in color for at least 30 seconds Table 4: Mass and Volume Titration Data for Equivalent Mass Determination of Maleic Acid Initial Final Initial Initial Final Final Trial mass mass HCI NaOH HCI NAOH (g) (g) (mL) (mL) (mL) (mL) 1 7.3022 7.2082 2.20 2.70 7.30 24.10 2 7.2082 7.1066 7.30 24.108.40 43.00 3 7.10666.8410 8.40 2.26 9.00 46.52 * Snapshot of data that you would have collected in the laboratory provided in RED
REPORT TABLE: Report the Equivalent Mass of the acid for all three runs, as well as the average value and all pertinent other values. ## Equivalent mass for maleic acid 58.035 g/mol (110-16-7. The Merck Index Online; Royal Society of Chemistry, 2013; M7037) ## Since we do not have regular access to the CRC or Merck Index as we typically would in our laboratory setting, the equivalent mass value was looked up and provided here with corresponding reference. As a reminder, always include the proper reference whenever making citations like this in your laboratory notebook/report.