Acid-Base Indicators and pH Chem 42 "Lab Fall 2020 Objectives Prepare solutions of known pH by dilution Observe colors o

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Acid-Base Indicators and pH Chem 42 "Lab Fall 2020 Objectives Prepare solutions of known pH by dilution Observe colors of various indicators from pH 1-13 Estimate the pH of household solutions from indicator colors Background Most students start chemistry with some idea of what pH is and that it is important in many practical situations. If you have ever maintained a hot tub or swimming pool, you probably have measured the pH of the water. You probably also know that pH 7 is considered neutral. You may know that below pH 7, a solution is considered acidic, while above pH 7 is considered basic. But what is pH? In short, pH reflects the concentration of hydronium (HO) ions in a solution. The more acidic (lower pH) a solution is the more of these hydronium ions are present. And the concentration of hydronium ions can be either quite high (as much as 1 mol/L or more in strongly acidic solutions) or quite low (below 10-** mol/L in very basic solutions). In most solutions, the concentration is less than 1 mol/L., and usually a lot less. See the table below for some examples: [H30* [OH-] 0.1 mol/L 10-13 mol/L 0.001 mol/L 10-11 mol/L 10-5 mol/L 10-mol/L 10-7 mol/L 10-7 mol/L 10-mol/L 10-5 mol/L 10-11 mol/L 0.001 mol/L 10-13 mol/L 0.1 mol/L PH 1 3 5 7 9 11 13 The pH of a solution is defined by the concentration of hydronium ions: pH = -log[H30) The logarithm is the power of the exponent in the concentration, and the negative means that most pH values will be positive. So neutral solutions are represented with a nice friendly number: pH 7 (rather than 1 x 10-7 mol/L H30“). Essentially. pH is a way of working with very small numbers in a more convenient way. What makes a solution acidic? In short, the more acids are dissolved in a solution, the more acidic it will be. Strong acids lower pH the most, but weak acids can lower it to a certain extent as well. A basic solution, of course, will have bases in it rather than acids. The most important ion in a basic solution is the hydroxide ion (OH-),
rather than the hydronium ion (although some hydronlum is always present). Why are these two particular ions important in water? They are the products of two water molecules reacting with each other (a process called autoionization): H20 (1) + H20 (1) H3O+ (aq) + OH- (aq) As a result, everything that contains water—your water bottle, your coffee, your oatmeal, your urine-will have both hydronium ions (acid) and hydroxide ions (base). Even the purest water will have both these ions, and one will always be present at a concentration of at least 10-7 mol/L. Granted, that is not very concentrated, but the concentration goes higher as the pH is farther from neutral. Recall that the concentration of the hydronium ions will determine the pH. (It is inversely related to hydroxide concentration as well). Method There are a number of ways to measure the pH of a solution. In this experiment we will look at two of the simplest: one is to use an indicator, a substance that changes color with pH. The second method is a pH meter, which is a cleverly designed electrode that gives different voltage readings at different pH values. Using a meter is relatively quick and accurate. First, solutions with known pH were prepared. You will study how they change the color of a number of indicators, and finally, you will use them to gauge the pH of some common household solutions. This data can then be compared to readings from a pH meter. Part I. Preparing solutions of known pH Solutions that are fairly acidic or basic can be prepared by diluting more concentrated solutions of acid or base. In this case, solutions of 0.1 mol/L HCl and 0.1 mol/L NaOH were used. Hydrochloric acid (HCI) is known as a strong acid because its ions separate completely in water to form H30* (hydronium) ions. HCI (aq) + H20 () → H30- (aq) + CH- (aq) This means that a 0.1 mol/L solution of HCl is really 0.1 mol/L of H3O+ ions. This solution is therefore pH 1. If the solution is diluted by a factor of 10, so that the concentration is 0.01 mol/L H307, then the pH will be 2. A third dilution to 0.001 mol/L results in a pH of 3. Notice that every change by one pH unit corresponds to a factor of 10 change in the concentration of H30* ions. Preparing solutions of pH 1, 2 and 3 can be done in this
way by using serial dilution. In a serial dilution, a solution is diluted, and then the diluted solution is diluted again. The concentration drops very rapidly as a result. Certain basic solutions can also be made by serial dilution. A 0.1 mol/L NaOH solution has a pH of 13, and diluting it by a factor of 10 changes the pH to 12 (less basic). A further dilution by a factor of 10 gives a pH of 11. Beyond this, we need to use a solution known as a buffer to have a stable pH. Watch the first video, which shows how diluting solutions of HCI (strong acid) and NaOH (strong base) corresponds to pH (see figure 1 for a diagram). Acid Solutions: pH 1: 10 mL of 0.1 mol/L HCl was added to the pH 1 test tube. pH 2: A plastic pipet was used to transfer 1 mL of the pH 1 solution into the pH 2 test tube. Then 9 ml of DI water was added to the pH 2 test tube and mixed. pH 3: 1 ml of the pH 2 solution was transferred to the pH 3 test tube, and 9 mL of DI water was added. Mixed well. Basic Solutions: pH 13:10 mL of 0.1 mol/L. NaOH was added to the test tube. pH 12: 1 mL of pH 13 solution and 9 ml. of Dl water were mixed in the pH 12 test tube pH 11:1 mL of pH 12 solution transferred to the pH 11 test tube and mixed with 9 ml. of Dl water. 10 ml 0.1 M 9 mL 9 ml 9 mL 9 mL 0.1 M HCI H2O H2O H2O H2O NaOH Buffers 10 mL 1 mL 1 mL Tut HUU 10 11 12 13 1 2 3 4 5 6 7 8 9 Figure 1. Preparation of pH 1-13 solutions
Part II. Testing Indicators at pH 1-13 Using both serial dilutions (from part 1) and buffers, solutions of pH 1-13 were made and transferred to a set of well plates for easy observation. A few drops of an indicator was added to each of the pH 1-13 solutions, and this was repeated for each of the indicators. Watch the second video (you can watch the colors change individually or speed to the end if you like). Each row corresponds to one acid-base indicator. In the table provided, record the color of each indicator at each pH value. Try to stick to descriptive names that are clear for most readers. pH-> 4 5 6 10 11 thymol blue (TB) thymolphthalein (T) methyl red (MR) bromocresol green (BG) bromothymol blue (BB) 1 2 3 7 8 9 12 13 Part 3. Finding the pH of Household Solutions Now that you have recorded the color each indicator at each pH, you will be able to estimate the pH of some common household solutions. Watch the third video and record the color of each indicator in each solution. Solution → Vinegar Baking Club soda Ammonia 3% soda cleaner hydrogen Indicator solution peroxide thymol blue (TB) thymolphthalein (T) methyl red (MR) bromocresol green (BG) bromothymol blue (BB) pH range indicators) pH meter reading
Using the colors you observed, estimate the pH of each solution and record it in the table. One quick way of doing this is to identify the pattern of five colors in the column of a particular solution. Then look for the same pattern of colors in the columns of the pH table you completed. For several of the solutions, this will point to very specific values for the pH. For others, there may be several possible pH values. Those should be reported as a range. Part IV. Checking with a pH meter. The pH of each solution was tested directly with a pH meter in the fourth video. Record the values that were measured in the table above. Note that it takes time for the meter to stabilize, and the readings just before the electrode is removed from the solution are the most accurate. Questions to Answer 1. How did the results from the pH meter compare with the pH you estimated using the indicators? Which solutions agreed well? Were there any that did not agree? 2. Using the dilution equation (MV= M2V2), calculate the concentration of NaOH in mol/L the pH 12 test tube. Notice that it was 0.1 mol/L at pH 13, and that 1 mL of this solution was added to 9 mL of Dl water to make the pH 12 solution. 3. Starting with the NaOH concentration in the pH 12 solution, now calculate the NaOH concentration in the pH 11 solution (again using the dilution equation). What do you notice? 4. Turn in pages 4 and 5 of this handout. This is the last Chem 42 lab! Thanks for your hard work and dedication in making it to this point in an unusual semester.