Name Team Name CHM101 Lab Acids and Bases Grading Rubric To participate in this lab you must have splash-proof goggles, proper shoes and attire. Criteria Points possible Points earned Lab Performance Printed lab handout and rubric was brought to lab 3 Safety and proper waste disposal procedures observed 2 Followed procedure correctly without depending too much on instructor or lab partner 3 Work space and glassware was cleaned up 1 Lab Report All observations accurately recorded 5 Question 1 1 Question 2 1 Question 3 1 Question 4 1 Question 5 2 Total 20 Subject to additional penalties at the discretion of the instructor.
Acids and Bases Goals: To determine the ph of common substances and observe buffer behavior. Background Acids and bases are very common in chemistry and biology. Understanding acids and bases is key to understanding many reactions in chemistry and biochemistry. The work acid comes from the Latin acidus meaning sour from acēre to be sharp. Acid solutions have a ph less than 7, turn blue litmus paper red, and taste sour. The name base was coined in the sense that bases react with acids and give them a solid form (turn them into salts). Bases have ph greater than 7, turn red litmus paper blue and tend to taste bitter. Acids and bases can be defined in different ways. The Arrhenius definition classified acids as substances which produce hydrogen ions (H + ) in water solutions (in water, H + really exists as hydronium (H 3O + ), but chemists often write just H + anyway). We can recognize them by the H in front of an anion with (aq). For example: HCl (aq) is an Arrhenius acid. Arrhenius bases are substances that dissociate into cations and hydroxide (OH - ) ions. For example: NaOH (aq) is an Arrhenius base. The Brønsted-Lowry definition expands the type of compounds that are considered acids and bases. A Brønsted-Lowry acid is defined as a proton (H + ) donor. By this definition, any compound that can lose H + can be considered an acid. Examples include all of the Arrhenius acids as well as compounds that we may not think of as acids like water and alcohols. A Brønsted Lowry base is defined as a substance that can accept H +. OH compounds will accept H + to make H 2O, but many other compounds will accept H + as well, such as ammonia (NH 3) or any compound with a lone pair. The Autoionization of Water and ph Water exists in equilibrium with hydronium and hydroxide (the autoionization of water.) H 2O(l) + H 2O(l) H 3O + (aq) + OH (aq) K w = [H 3O + ][OH ] The equilibrium constant for this reaction, K w, is equal to the hydronium time hydroxide concentration. At 25 o C, the value of K w is 1.0 x 10 14. In pure water at 25 ⁰C, [H 3O + ] = 1.0 x 10 7 M and [OH ] = 1.0 x 10 7 M. The K w formula works for non-neutral solutions also. If the [H 3O + ] > [OH ], the solution is acidic and if [OH ] > [H 3O + ] then the solution is basic. ph is a way to express the acidity or basicity of a solution without having to write out concentrations with exponents.. ph = - log[h 3O + ] In neutral water ph = -log [1.0 x 10 7 M] = 7. Acidic solutions have ph less than 7. Basic solutions have ph greater than 7.
Here are some example values for [H + ] and ph, and next to them, what [OH ] and poh will be in the same solution. [H 3O + ] (molarity) ph [OH ] (molarity) poh 1 x 10 1 = 10 M -1 1 x 10 15 M 15 1 x 10 0 = 1 0 1 x 10 14 14 1 x 10 1 = 0.1 1 1 x 10 13 13 1 x 10 2 = 0.01 2 1 x 10 12 12 1 x 10 3 = 0.001 3 1 x 10 11 11 1 x 10 4 = 0.0001 4 1 x 10 10 10 1 x 10 5 = 0.00001 5 1 x 10 9 9 1 x 10 6 = 0.000001 6 1 x 10 8 8 1 x 10 7 = 0.0000001 7 1 x 10 7 7 1 x 10 8 = 0.00000001 8 1 x 10 6 6 1 x 10 9 = 0.000000001 9 1 x 10 5 5 1 x 10 10 = 0.0000000001 10 1 x 10 4 4 1 x 10 11 = 0.00000000001 11 1 x 10 3 3 1 x 10 12 = 0.000000000001 12 1 x 10 2 2 1 x 10 13 = 0.0000000000001 13 1 x 10 1 1 1 x 10 14 = 0.00000000000001 14 1 x 10 0 0 1 x 10 15 = 0.000000000000001 15 1 x 10 1-1 There are many ways to measure ph. One such way is to use an indicator, which is a compound that changes color as the ph changes. Red litmus paper turns blue in the presence of a base, for example. A universal indicator contains several individual indicators mixed together the combined colors that can be used to estimate the ph in the range 1-14. In this lab, we will use universal indicator solution, ph paper, or a red cabbage solution to estimate ph. Indicators are usually accurate to within one ph unit. To get more accurate measurements of ph, we will use a ph meter. The ph meter works by carefully measuring the voltage difference between your test solution, and another solution inside the probe with ph = 7. More ions in your solution creates a higher voltage difference. The probe uses a special glass barrier that only responds to H +, so that it will measure H + alone, without other ions interfering. This way, the ph can be measured to a few decimal places, especially with a high-quality ph meter. Buffers A buffer solution is a solution that resists a change in its ph upon the addition of small quantities of either a strong acid or a strong base. Buffers are usually made by mixing a weak acid and its conjugate base, or a weak base and its conjugate acid. For example, a solution containing NH 4Cl with NH 3 will be a buffer solution. In this lab we will make a buffer using HC 2H 3O 2 (aq) and NaC 2H 3O 2 (aq). Buffers work by neutralizing the added strong acid or base. When a strong acid is added (H 3O + ), it reacts with the weak base (C 2H 3O 2_ ) creating weak acid (HC 2H 3O 2) and water. Since the strong acid has reacted, the ph stays relatively stable. C 2H 3O 2 _ (aq) + H 3O + (aq) HC 2H 3O 2 (aq) + H 2O When a strong base (OH ) is added to a buffer it will react with the weak acid (HC 2H 3O 2) and produce a weak base (C 2H 3O 2_ ) and water. Since the strong base has reacted, the ph stays relatively stable. HC 2H 3O 2 (aq) + OH (aq) C 2H 3O 2 _ (aq) + H 2O
Laboratory Activity Materials: premade universal indicator solutions, ph 1-14 various household substances stirring rod universal indicator solution ph paper ph meter 0.10 M HCl 0.10 M NaOH Buffer solution (0.10 M sodium acetate in 0.10M acetic acid) Procedure A. Determination of ph of Household Substances 1. Your instructor will prepare a series of test tubes containing solutions of ph 1-13 treated with universal indicator. Record the color of each tube. Do not open the tubes! 2. Obtain 10 household substances. If the substance is a liquid, place 1-2 ml in a test tube. If the substance is a solid, add one spatula tip to the test tube then add about 1 ml of water. If the solid is in chunks or pill form, grind with a mortar and pestle before adding to the test tube. 3. Determine the ph using ph paper. To use ph paper: Do not dip the paper itself in the solution the indicators will dissolve in the solution. Dip a clean, dry stirring rod in the test tube. Remove the rod and touch it to ph paper. Estimate the ph by comparing the color of the paper to the color chart on the package 4. Determine the ph using universal indicator solution. To use universal indicator solution: Add 2 drops of indictor to the test tube and shake to distribute. Estimate the ph by comparing to the prepared solutions in step 1. B. ph of Unbuffered and Buffered Solutions 1. Clean two 25 ml beakers and place 10.0 ml of deionized water in each. Use a graduated cylinder to measure the required volume of water. 2. It is difficult to determine the ph of deionized water, so assume that the ph = 7. 3. Measure 1 ml of 0.10 M HCl with a pipette and add to one beaker. Stir and measure the ph with a ph meter. To use a ph meter: Rinse the probe with deionized water. Completely submerge the glass sphere of the probe in the liquid. Rinse again with deionized water after each reading. 4. Measure 1 ml of 0.10 M NaOH with a pipette and add to the other beaker. Stir and measure the ph meter. 5. Clean two new 25 ml beakers and place 10.0 ml of buffer solution (0.10M HC 2H 3O 2, 0.10 M NaC 2H 3O 2) in each. 6. Measure the ph of the buffer solution before adding acid and base. 7. Measure 1 ml of 0.10 M HCl with a pipette and add to one beaker of buffer. Stir and measure the ph meter. 8. Measure 1 ml of 0.10 M NaOH with a pipette and add to the other beaker of buffer. Stir and measure the ph meter. Waste Disposal Pour all solutions down the sink. Be sure to rinse all solid from test tubes before putting back.
Acids and Bases: Data Sheet Name A. Determination of ph of Household Substances Prepared Universal Indicator Solutions: ph Color Observed ph Color Observed 8 1 9 2 10 3 11 4 12 5 13 6 7 Substance ph with indicator paper ph with universal indicator acidic, basic or neutral? Q1. Which household substance was the most acidic? Which household substance was the most basic? Report Page 1 of 3
Acids and Bases: Post Lab Name B. ph of Unbuffered and Buffered Solutions Solution ph Deionized water 7 Deionized water + 1 ml 0.10 M HCl Deionized water + 1 ml 0.10 M NaOH Buffer (0.10M HC 2H 3O 2, 0.10 M NaC 2H 3O 2) Buffer + 1 ml 0.10 M HCl Buffer + 1 ml 0.10 M NaOH Q2. a) Describe the difference in behavior of pure water and buffer solution when strong acids and bases are added: b) Why do buffers behave differently? Q3. Based on your measured ph, calculate the hydronium concentration of the solution: Deionized water + 2.5 ml 0.10 M NaOH. (show work) Q4. Based on your measured ph, calculate the hydroxide concentration of the solution: Buffer + 2.5 ml 0.10 M HCl. (show work) Report Page 3 of 3
Acids and Bases: Post Lab Name Q5. Complete the following table: [H 3O + ] [OH - ] ph acidic, basic, or neutral? 1.66 x 10-5 M 6.88 x 10-2 M 12.26 5.01 x 10-13 M Report Page 3 of 3