Bio 3A LAB: Acids and Bases (rev.2/3/2009) Page 1 of 4 Biology 3A Laboratory Acids, Bases and Buffers Objectives • To understand the concept of pH • To be able to calculate pH from acid molar concentration • To measure pH using instrumentation and indicators • To understanding the effect of a buffer on pH Introduction This lab assumes a knowledge of the chemical concept of molar concentration. An important application of molarity, or molar concentration, involves hydrogen ions in solution. The concentration of free hydrogen ions in pure water is 1.0 x 10 -7 M. This occurs because water can dissociate as shown in this equation (1) In pure water, the concentration of OH- and H+ ions are equal (both are 1 x 10 -7 M). If we increase the concentration of hydrogen ions the solution becomes acidic; if we decrease the concentration of hydrogen ions the solution will become basic. Chemists use the pH scale to indicate acidity. In this scale, pH is equal to the negative logarithm of the hydrogen ion concentration pH = - log [H+] (2) The pH scale generally ranges from 0.0 to 14.0. Pure water has a pH of 7.0. If we consider 1.0 M HCl, it has a hydrogen ion concentration of 1 x 100. Thus it has a pH of 0.0. If we were to have a 10.0M HCl solution, its pH would be -1.0! In this lab we will look the problem of measuring pH using indicators and instruments, and then consider ways to control the fluctuation of pH. In this lab we will use the notation M to indicate molar solutions (number of moles per liter) and mM to indicate millimolar solutions (number of millimoles per liter). If a solution has 1 mole per liter, that is also 1000 millimoles per liter; therefore 1 M = 1000 mM. pH Indicators An indicator is a chemical that changes predictably under changing condition. There are several well-known pH indicators that change color at specific levels of hydrogen ion concentration. Five of the well-known pH indicators are shown in Table 1. Indicator pH Color in Acid Color in Base Phenolphthalein 8.3 – 10.0 Colorless Red - Pink Litmus 4.5 – 8.3 Red Blue Bromcresol Purple 5.2 – 6.8 Yellow Purple Bromophenol Blue 3.0 – 4.6 Yellow Blue Methyl Violet 0.2 – 3.0 Yellow Blue - Violet Table 1. Some common pH indicator chemicals and their respective colors in acidic and basic solutions Bio 3A LAB: Acids and Bases (rev.2/3/2009) Page 2 of 4 Litmus, a common biological pH indicator In Table 1, litmus is a pH indicator that seems to broadly cover the neutral range of pH (both sides of 7.0). Litmus is a naturally occurring compound found in many plants. If you are familiar with the Hydrangea, you may know that it can occur in both a blue and a pink form. Actually, it’s the same plant and same flowers; it’s just the pH of the soil that alters the color! Grow it in acidic soil and it’s pink; grow it in basic soil and it’s blue. Another source for litmus is red cabbage (or perhaps blue cabbage, depending on the pH). We can use red cabbage to create a litmus solution, which we can then use to observe pH changes. Procedure 1. Make and use Litmus from red cabbage 1. One group will make enough for the entire class by cutting a handful of red cabbage leaves. Using the blender and a small amount of tap water, shred the leaves. 2. Strain the mixture through a double layer of cheesecloth 3. Decant about about 30 ml of the solution into a beaker. This is the litmus solution. What color is it right now? What is the current pH (use pH paper)? 4. Add about 80 mL tap water. What can you conclude about the pH tap water here Saddleback College? 5. Add 0.1 M HCl drop by drop until you see a color change. Describe the color change. What is the current pH of the solution? If you were going to make a cabbage salad and you wanted the red cabbage to be bright red, what dressing could you use? (If this question is unclear, answer it after you complete the following section.) Procedure 2. Measuring the pH of common solutions Some indicators have a much broader pH range. Hydrion® paper is an example of a broad range indicator conveniently applied to a paper used for pH testing. Using Hydrion® pH paper, measure and record pH of some common solutions provided in lab. Enter the measured pH values into Table 2. Procedure 3. Using a pH meter A pH meter can be used to directly measure the hydrogen ion concentration. Please watch your instructor’s demonstration of the pH meter. Always wash the electrode with deionized water after use. Always place the washed electrode into deionized water after washing. In procedure 2 you measured the pH of three HCl solutions. Use the meter to measure the pH of each of these solutions again. Enter your measurements into Table 2. Procedure 4. Compare Calculated pH with pH paper and pH meter Let’s calculate the actual free hydrogen ion concentration in an acid solution of known concentration. In the earlier exercise you measured the pH of a 0.01 M HCl solution. In an aqueous solution HCl completely dissociates, or in other words, the H+ and the Cl- ion part company totally. Thus if we known the concentration of HCl, that is equal to the free hydrogen ion concentration or [H+]. So in our 0.01 M HCl we have a free hydrogen ion concentration of 1.0 x 10-2 moles per liter. pH = - log [H+] = - log [1.0 x 10-2] = 2. So the pH of 0.01 M HCl should be 2.0Bio 3A LAB: Acids and Bases (rev.2/3/2009) Page 3 of 4 Calculate the expected pH for 0.001 M HCl and 0.0001 M HCl. Enter these expected values into Table 2. Compare your measured values using pH paper and the pH meter to the calculated values. Please discuss any variability in your results from each of the three measurements. Procedure 5. How effective are antacids? Typical over-the-counter (OTC) antacids claim to “neutralize excess stomach acid.” You’ve seen the commercials and advertisements. Since the acid produced by your stomach is HCl, we can easily find out how well they work. The typical pH of the acid in your stomach is around 2.0. In lab you will find a selection of OTC antacids. Test each of these in the following manner. In addition,