Acids, Bases, Buffers and Salts-3 PDF

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Acids, Bases, Buffers and Salts Experiment #3 Week#: 3

CHEM 114.1- General Chemistry II Redesigned Laboratory

Queens College October 29th, 2019

Abstract:

Within this particular experiment, it was determined that the concentration of an acid, which was controllable, and the strength of an acid, which was out of control, both took part as to whether or not a reaction will result in an acidic or basic solution. Hydronium ions result in an acidic solution, while hydroxide ions result in a basic solution. Through a reaction with water, and different reactions, various results take place, in which using strong acid and weak acid may result in similar pH depending on the concentration of each compound used. as opposed to ammonia. A reaction with buffer and water showed a resistance in change of pH, while a presence of a strong acid, such as hydrochloric acid caused a dramatic change depending on the concentration values. Salt were determined do react differently with water, based off of the pH indicator strips, which provided information as to which products have formed depending on the solutions basicity or acidity.

Introduction: For purposes throughout the lab, an acid is defined as a substance that increases the concentration of hydrogen ions (H+) in a solution. This usually happens as the acid donates one of its hydrogen atoms by means of dissociation. A base, on the other hand, raises the pH of a system by increasing the concentration of hydroxide (OH_). Bases act as hydrogen acceptors and act in conjugation with acids as well. With regards to the relative strengths of acids, the stronger acids will more readily dissolve in water and release hydrogen ions. Similarly, strong bases also completely dissociate in water and release hydroxide ions to combine with hydrogen ions. A salt is defined as an ionic compound composed of two groups of oppositely charged compounds. Salts can be easily identified since they usually consist of positive ions from a metal with negative ions from a nonmetal. Buffers, on the other hand are solutions that can

resist changes in pH when acids and bases are added to them. The added acids/bases come in the form of OH- or H3O+ ions, and are neutralized if the buffer is strong enough to do so. Another crucial component of understanding buffers is buffer capacity. Buffer capacity refers to the amount of acid or base that can be added to a buffer before its pH shifts by 1 unit from its equilibrium position. This capacity varies from buffer to buffer and is dependent upon the volume and concentrations of the acid and base that compose it. This lab will utilize acid-base indicators to test solutions for the presence of one of the two components. Each solution will react differently with each indicator, allowing one to draw conclusions from clear observations.Concentrations acid and bases can be determined using the equation M1V1 = M2V2, which can be set up in the following units. (Molarity of Previous Test Tube) (Volume transferred mL) = (New Molarity) (Desired Volume in mL)

Experimental: In the first part of the experiment, 3 small beakers are obtained and are labeled 1, 2, and 3. 50 mL of 0/10 M HCl was poured in the first beaker. 5.0 mL of the solution was poured into a 100 mL graduated cylinder and distilled water was added until the 50.0 mL mark. The mixture was then poured into the second beaker and was mixed well. 5.0 mL of the solution in the second beaker was poured into the graduated cylinder and was diluted to 50.0 mL. The mixture was poured into the third beaker and was mixed well. The pH of each solution was measured with a pH meter. The electrode from the solution was removed the solution and was rinsed with distilled water. The electrode was placed in the solution to be measured. The bulb of the pH electrode was made sure to be submerged in the solution. The value of the pH displayed was recorded. The electrode was then removed from the solution and was prepared for the next two

solutions. The beakers were rinsed and dried, and the procedure was repeated using 0.10 M NaOH. The experiment was repeated using 1.0 M HC2H3O2 and 1.0 M NH3. The concentrations of HC2H3O2 or NH3 in each of the diluted solutions were calculated. The pH measurement made for each HC2H3O2 solution was used to calculate a value of K a  and pK a  for HC2H3O2. The pH measurement made for each NH3 solution was used to calculate a value of K b  a  nd pK a  for NH3. In the second part of the experiment, three serial dilutions are prepared to the buffer solution made to contain 1.0 M acetic acid and 1.0 M sodium acetate. 5.0 mL of solution from the third beaker was removed so that all three beakers contain the same volume if buffer solution. The pH of each solution was measured. 10.0 mL of 0.10 M HCl was added to each solution and mixed well. The pH of the challenged buffer solutions was measured. The change in pH produced by the addition of HCl to each buffer solution was calculated. In the third part of the experiment, a spatula tip full for each of the ionic compounds was placed in a small sized test tube half filled with distilled water. The solution was mixed well to dissolve the solid. A small sample was removed from the solution and was touched to a piece of pH paper. The color of the paper was compared to the chart on the side of the container to estimate the pH of the solution to the nearest pH unit.

Discussion: In part 1 of the experiment the pH values of the 3 diluted solutions were recorded along with the concentrations of HCl or NaOH. Solution 1 obtained a pH of 1.25, with an HCl concentration of 0.1 M. Solution 2 obtained a pH of 2.50 with an HCl concentration of 0.01 M. Solution 3 obtained a pH of 3.43 with an HCl concentration of 0.001 M. In the repeated procedure using NaOH instead of HCl, solution 1 obtained a pH of 11.21 with an NaOH concentration of 0.1 M. Solution 2 obtained a pH of 10.54 with an NaOH concentration of 0.01 M. Solution 3 obtained a pH of 9.88 with an NaOH concentration of 0.001 M. The pH values were then measured from diluted solutions of 1.0 M HC2H3O2 and 1.0 M NH3. In solution 1, the pH was recorded to be 2.25 and the concentration of HC2H3O2 was 1.0 M. In solution 2, the pH was recorded to be 3.51 and the concentration of HC2H3O2 was 0.1 M. In solution 3, the pH was recorded to be 3.75 and the concentration of HC2H3O2 was 0.01 M.  and the pKa  was calculated to be For solution 1 the Ka  value was calculated to be 3.18 x 10-5  and the pKa  was calculated to 4.50. For solution 2 the Ka  value was calculated to be 9.58 x 10-7  and the pKa  was calculated be 6.02. For solution 3 the Ka  value was calculated to be 3.22 x 10-6 to be 5.49. All three of the pKa  values were similar. Calculation of Ka  and pKa  (HC2H3O2 solution): Solution 1: Ka=pKa = -log(3.18 x 10-5  ) = 4.50 Solution 2: Ka=pKa = -log(.58 x 10-7  ) = 6.02

Solution 3: Ka= pKa = -log(3.22 x 10-6  ) = 5.49 The procedure was repeated by using NH3 instead of HC2H3O2. Solution 1 obtained a pH of 10.55 with an NH3 concentration of 1.0 M. Solution 2 obtained a pH of 9.98 with an NH3 concentration of 0.1 M. Solution 3 obtained a pH of 9.39 with an NH3 concentration of 0.001M. Your friend tells you that the pH of a solution of a strong acid must be lower than the pH of a solution of a weak acid. This statement would be incorrect because the pH of an acid depends on its concentration of hydrogen ions which can be changed by dilution. The pH measurement of 0.1 M NaOH was 11.21 and for 0.1 M NH3 was 9.89. These measurements show that NaOH is a stronger base than NH3 because NaOH has a stronger base than NH3. For solution 1 the Kb value was calculated to be and the pKb  was calculated to be 6.90. For solution 2 the Kb  value was calculated to be and the pKb  was calculated to be 7.22. For solution 3 the Kb  value was calculated to be and the pKb  was calculated to be 7.22. All three pKb  values were similar.

Calculation of pOH (NH3 solution) Solution 1: pOH = 14 - 10.55 = 3.45 Solution 2: pOH = 14 - 9.89 = 4.11 Solution 3: pOH = 14 - 9.39 = 4.61

Calculation of Kb  and pKb  (NH3 solution): Solution 1:

Kb= pKb = -log(1.26 x 10-7  ) = 6.90 Solution 2: Kb=pKb = -log(6.03 x 10-8  ) = 7.22 Solution 3: Kb= pKb = -log(6.04 x 10-8  ) = 7.22 In part 2 of the experiment, the pH values were measured from buffer solutions. In solution 1 the pH was 4.72, solution 2 had a pH of 4.59, and solution 3 had a pH of 4.40. When HCl was added to each solution, their pH decreased. The pH of solution 1 became 4.32, the pH of solution 2 became 4.11, and the pH of solution 3 became 2.26. The change in pH for solution 1 was 0.40, 0.48 in solution 2, and 2.14 in solution 3. In part 3, Na2CO3 had a dark green color and a pH of 10. NH4 had a green color and a pH of 5. (NH4)2CO3 had a yellow/orange color with a pH of 9. AlCl3 had a dark orange color and a pH of 3. Based on the pH measured for the solution of (NH4)2CO3, CO3 Chemical reactions: CO32- + H2O ⇆ HCO3- + OHNH4 + H2O ⇆ NH3 + H3O+ NH4 + H2O ⇆ NH3 + H3O+ 3Cl- + H2O ⇆ HCl + H3O+

2-

is a stronger base.

Conclusion: The goal of this experiment was to learn about acids, bases, and salts, along with their chemical properties. We tested different things such as pH, solubility, and how to interpret and observe chemical reactions so that we could then write out their equations. Almost all chemical compounds can be identified as an acid, base, or salt. By understanding the nomenclature of these compounds, we can express chemical reactions and balance these equations to better understand their stoichiometry. Two errors that could have occurred during the experiment was that I mixed up the tubes and their solutions resulting in a mix up of the pH readings. A second error that could have occurred could have been that I didn’t use the correct volumes of the solutions in the tubes. Perhaps this could alter the pH reading due to the difference in volume

SWI: After completing this experiment I clearly understood acids, bases, buffers and salts along with their chemical properties. W tested different things about acids, bases and salts, along with their I think additional guidance was needed when creating the solutions as my partner and I were a bit confused with some steps. Throughout the experiment I think the instructions could be more straight to the point. Along with that, I noticed it was easier to use the lab quest, since we have been using it over the past few weeks. Overall, I think it was a straight forward, and enjoyable lab....