Acids and Bases Lab Procedure PDF

# Acids, Bases, pH, and Buffers ## Laboratory Goals - Extract a naturally occurring dye from red cabbage to use as a pH indicator. - Measure the pH of several substances using the cabbage indicator and a pH meter. - Calculate pH from the [H₃O⁺] or the [OH^-] of a sol...

# Acids, Bases, pH, and Buffers ## Laboratory Goals - Extract a naturally occurring dye from red cabbage to use as a pH indicator. - Measure the pH of several substances using the cabbage indicator and a pH meter. - Calculate pH from the [H₃O⁺] or the [OH^-] of a solution. - Calculate the molar concentration and percentage of acetic acid in vinegar. - Observe the changes in pH as acid or base is added to buffered and unbuffered solutions. ## Lab Information - **Time:** 2 ½ h - **Comments:** - Students may be asked to bring a red cabbage and/or colorless household samples to class. - Share test tubes with your lab neighbors to prepare the pH reference solutions. - Tear out the report sheets and place them beside the matching procedures. ## Related Topics - Acids, bases, pH, buffers ## Chemical Concepts - *An acid is a substance that dissolves in water and donates a hydrogen ion, or proton (H⁺), to water.* In the laboratory, we have been using acids such as hydrochloric acid (HCl) and nitric acid (HNO₃). - HCl + H₂O --> H₃O⁺ + Cl^- - You use acids and bases every day. For example, there are acids in oranges, lemons, and vinegar. In this experiment we will use acetic acid (HC₂H₃O₂). Acetic acid is the acid in vinegar that gives it a sour taste. - Some typical bases used in the laboratory are sodium hydroxide (NaOH) and potassium hydroxide (KOH). Most of the common bases dissolve in water and produce hydroxide ions, OH^-. - A common base produces a cation and an OH^- anion in an aqueous solution. - An important weak base found in the laboratory and in some household cleaners is ammonia. In water, ammonia accepts H⁺ to form ammonium and hydroxide ions. - NH₃ + H₂O ---> NH₄⁺ + OH^- ## A. Reference Colors for pH Using Red Cabbage Indicator - The pH of a solution tells us whether a solution is acidic, basic, or neutral. On the pH scale, values below 7 are acidic, a value equal to 7 is neutral, and values above 7 are basic. Typically, the pH scale has values between 0 and 14. ## B. Measuring pH - The molar concentrations of H₃O⁺ and OH^- are indicated by square brackets. In pure water, [H₃O⁺] = [OH^-]=1.0×10^-7 M at 25°C. The ion constant product of water, Kw, can thus be calculated. - Kw = [H₃O⁺][OH^-]=[1.0×10^-7][1.0×10^-7]=1.0×10^-14 - If the [H₃O⁺] or [OH^-] for an acid or a base is known, the other can be calculated. For example, if an acid has a [H₃O⁺]=1.0×10^-4 M, we can find the [OH^-] of the solution by solving the Kw expression for [OH^-]. - [OH^-] = 1.0×10^-14/ [H₃O⁺] = 1.0×10^-14/1.0×10^-4 =1.0×10^-10 M - The pH of a solution is a measure of its [H₃O⁺]. It is defined as the negative log of the hydrogen ion concentration. - pH = -log [H₃O⁺] - Therefore, a solution with [H₃O⁺]=1.0×10^-9 M has a pH of 9.00 and is basic. - pH = -log [1.0×10^-9] - pH = -[-9.00] - pH = 9.00 - The number of significant figures to the right of the decimal point in the pH is equal to the number of significant figures in the coefficient of the [H₃O⁺]. ## C. Effect of Buffers on pH - The pH of the blood is maintained between 7.35 and 7.45 by buffers in the body. If blood pH goes above or below that range, it can damage or destroy the cells in the blood. Buffers maintain the pH of a solution by reacting with and neutralizing small amounts of acids or bases. Many buffers contain a weak acid and its salt. The weak acid reacts with excess base, and the anion of the salt picks up excess H⁺. It is the ability of a buffer to react with excess acid or base that maintains the pH of a solution. One important buffer system in the blood is the hydrogen carbonate buffer, which is carbonic acid, H₂CO₃ (weak acid) and hydrogen carbonate anion, HCO₃^- (salt). When base (OH^-) is added, it reacts with the weak acid in the buffer and produces hydrogen carbonate ion and water. - H₂CO₃(aq) + OH^- (aq) → HCO₃^-(aq) + H₂O(l) - When acid (H⁺) enters the blood, it reacts with the HCO₃^- anion and re-forms carbonic acid: - H₃O⁺(aq) + HCO₃^-(aq) → H₂CO₃(aq) + H₂O(l) ## Experimental Procedures ### A. Reference Colors for pH Using Red Cabbage Indicator - **Materials:** Red cabbage leaves; 400-mL beaker; distilled water; Bunsen burner or hot plate; small (100-mL or 150-mL) beaker; 13 test tubes; two test tube racks; set of buffers with pH ranging from 1 to 13 #### Preparing red cabbage indicator - Tightly pack torn red cabbage leaves in a 150-mL beaker, then transfer leaves to a 400-mL beaker. Add about 150 mL of distilled water or enough water to cover the leaves. Heat, using a Bunsen burner or a hot plate, but do not boil. When the solution has attained a dark purple color, turn off the heat and allow the solution to cool. #### Preparing pH reference standards - Place about 30 mL of cabbage dye indicator in a small beaker. Arrange 13 labeled test tubes in two test tube racks. You may need to combine your test tubes with those of your neighbor (your instructor may prepare a pH reference set for the entire class). Pour 2 mL of each buffer solution in separate test tubes to give a pH reference set with pH 1-13. **Caution: Low pH values are strongly acidic; high pH values are strongly basic. Work with care. To each test tube, add about 2 mL of the cooled red cabbage solution. To obtain a deeper color, add more cabbage solution. Shake the test tube to mix. Record the color of the each of the solutions at different pH values. Keep this reference set for the next part of the experiment.** ### B. Measuring pH - **Materials:** Shell vials or test tubes; samples to test for pH (shampoo, conditioner, mouthwash, antacids, detergents, fruit juice, vinegar, cleaners, aspirin, etc.); cabbage juice indicator from part A; pH meter; calibration buffers; wash bottle; Kimwipes. #### Using red cabbage indicator to determine pH 1. Place 2 mL of each liquid sample in a shell vial (or a test tube). Add 2 mL of red cabbage solution to each sample. Shake the test tube to mix. Record the color. Prepare solutions of solid/viscous samples by adding small amount to 2 mL of water and mix thoroughly. 2. Compare the color obtained is part 1 to the colors of the pH reference set. Identify and record the pH that gives the best color match to the color of your sample. Repeat for other samples. **Dispose of all chemicals as directed by your lab instructor.** #### Using a pH meter to determine pH 3. **Your instructor will demonstrate the use of the pH meter.** 4. Using a new sample of a solution, determine the pH with the pH meter. Rinse off the pH electrode and use the pH meter to find the pH of other samples. 5. Identify each sample as acidic, basic, or neutral. ### C. Effect of Buffers on pH - **Materials:** Buffer with a high pH (9-11), buffer with a low pH (3-4), droppers, test tubes, graduated cylinder, 0.1 M NaCl, 0.1 M HCl, 0.1 M NaOH, pH meter, cabbage juice indicator from part A. - **Use a graduated cylinder to measure each of the following solutions and place in a separate test tube:** - 10 mL of H₂O - 10 mL of a buffer with a high pH - 10 mL of 0.1 M NaCl - 10 mL of a buffer with a low pH #### C.1 Effect of adding acid 1. Determine the pH of each by adding 2 mL of cabbage indicator to each sample and or by using a pH meter. Record. 2. Add 5 drops of 0.1 M HCl solution to each of the four test tubes from part 1 and mix. Determine the pH of each using your pH reference solutions or pH meter. Record. 3. Add 5 more drops of 0.1 M HCl solution to each of the four test tubes from part 2 and mix. Determine the pH of each using your pH reference solutions or pH meter. Record. 4. Determine the pH change in each of the solutions. 5. Identify the solutions that are buffers. #### C.2 Effect of adding base 1. Determine the pH of each by adding 2 mL of cabbage indicator to each sample or by using a pH meter. Record. 2. Add 5 drops of 0.1 M NaOH solution to each of the four test tubes from part 1 and mix. Determine the pH of each using your pH reference solutions or pH meter. Record. 3. Add 5 more drops of 0.1 M NaOH solution to each of the four test tubes from part 2 and mix. Determine the pH of each using your pH reference solutions or pH meter. Record. 4. Determine the pH change in each of the solutions. 5. Identify the solutions that are buffers. ## Pre-Lab Study Questions 1. How is the pH of a solution related to the [H₃O⁺]? 2. Using the equation for K_w, explain how [OH^-] changes when more H₃O⁺ is added. 3. Is a solution with a pH of 12.0 acidic or basic? 4. Is a solution with a pH of 2.0 acidic or basic? 5. What is a buffer? 6. If you add acid or base to water, how will the pH change? 7. If you add acid or base to a buffer, how will the pH change? ## Report Sheet ### A. Reference Colors for pH Using Red Cabbage Indicator | pH | Colors of Acidic Solutions | pH | Colors of Basic Solutions | |---|---|---|---| | 1 | | 8 | | | 2 | | 9 | | | 3 | | 10 | | | 4 | | 11 | | | 5 | | 12 | | | 6 | | 13 | | | 7 | Color of Neutral Solution | | | ### B. Measuring pH | Substance | 1. Color with Indicator | 2. pH Using Indicator | 3. pH Using pH Meter | 4. Acidic, Basic, or Neutral? | |---|---|---|---|---| | Household cleaners | | | | | | Vinegar | | | | | | Ammonia | | | | | | Drinks, juices | | | | | | Lemon juice | | | | | | Apple juice | | | | | | Detergents, shampoos | | | | | | Shampoo | | | | | | Detergent | | | | | | Hair conditioner | | | | | | Health aids | | | | | | Mouthwash | | | | | | Antacid | | | | | | Aspirin | | | | | | Other items | | | | | ### Questions and Problems **Q1** Complete the following table: |[H₃O⁺]| [OH^-]| pH | Acidic, Basic, or Neutral? | |---|---|---|---| | 2.0×10^-6| | 9.8 | | | | 3.5×10^-3 | | Neutral | **Q2** A solution has a [OH^-]=4.0×10^-5 M. What are the [H₃O⁺] and the pH of the solution? **Q3** A sample of 0.0084 mol of HCl is dissolved in water to make a 1500-mL solution. Calculate the molarity of the HCl solution, the [H₃O⁺], and the pH. For a strong acid such as HCl, the [H₃O⁺] is the same as the molarity of the HCl solution. HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl^-(aq) ### C. Effect of Buffers on pH #### C.1 Effect of adding acid | Solution | 1. Initial pH | 2. pH after 5 drops of HCl | 3. pH after 10 drops of HCl | 4. pH change | 5. Buffer yes or no? | |---|---|---|---|---|---| | H₂O | | | | | | | 0.1 M NaCl | | | | | | | High pH buffer | | | | | | | Low pH buffer | | | | | | #### C.2 Effect of adding base | Solution | 1. Initial pH | 2. pH after 5 drops of NaOH | 3. pH after 10 drops of NaOH | 4. pH change | 5. Buffer yes or no? | |---|---|---|---|---|---| | H₂O | | | | | | | 0.1 M NaCl | | | | | | | High pH buffer | | | | | | | Low pH buffer | | | | | | ### Questions and Problems **Q4** Which solution(s) showed the greatest change in pH? Why? **Q5** Which solutions(s) showed little or no change in pH? Why? **Q6** Normally, the pH of the human body is fixed in a very narrow range between 7.35 and 7.45. A patient with an acidotic blood pH of 7.3 may be treated with an alkali such as sodium hydrogen carbonate. Why would this treatment raise the pH of the blood?