Acids, Bases and pH Theory: An acid (from the Latin “acidus”, meaning “sour”) is a substance which, in water solution, has a sour taste, turns blue litmus red, and reacts with bases and certain metals (like magnesium) to form salts. Aqueous solutions of acids have a pH of less than 7. A lower pH means higher acidity, and thus a higher concentration of H+ ions (protons) in the solution. There are three common definitions for acids: the Arrhenius definition, the Brønsted-Lowry definition, and the Lewis definition. The Arrhenius definition defines acids as substances which produce hydrogen ions (H +) in water solutions (in water, H+ really exists as hydronium (H3O+), but chemists often write just H+ anyway). The Brønsted-Lowry definition expands on this: an acid is a proton (H+) donor. By this definition, any compound which can easily be deprotonated (lose H+) can be considered an acid. Examples include all of the Arrhenius acids as well as alcohols and amines which contain O-H or N-H fragments. The Lewis definition expands this further: a Lewis acid is a substance that can accept a pair of electrons to form a covalent bond. H+ does exactly that when it attaches to bases (it accepts an electron pair from the base to bond with it), but some other chemicals do that too. Examples of Lewis acids include H+, all metal cations, and electrondeficient molecules such as boron trifluoride and aluminium trichloride. Common examples of acids include hydrochloric acid (a solution of hydrogen chloride which is found in gastric acid in the stomach and activates digestive enzymes), acetic acid (vinegar is a dilute solution of this liquid in water), sulfuric acid (used in car batteries), and tartaric acid (a solid used in baking). As these examples show, acids can be solutions or pure substances, and can be derived from solids, liquids, or gases. Strong acids and some concentrated weak acids are often dangerous and corrosive. In chemistry, a base is a substance that, in water solution, is slippery to the touch, tastes bitter, turns red litmus paper blue, reacts with acids to form salts, and speeds up certain chemical reactions (bases can be catalysts). According to Arrhenius, bases produce hydroxide ions (OH−) in water solution. Examples of Arrhenius bases are the hydroxide compounds of the alkali and alkaline earth metals (NaOH, KOH, Ca(OH)2, etc.). Like the acid definitions, Brønsted–Lowry expands on Arrhenius, and Lewis expands on Brønsted–Lowry. A Brønsted–Lowry base is a substance that can accept H+. OH− compounds will accept H+ to make H2O, but many other compounds will accept H+ as well, such as ammonia (NH3) and amines (or any compound with a lone pair on a nitrogen (N:) ). A Lewis base donates an electron pair to form a covalent bond, but since accepting H+ means donating an electron pair to bond with the H+, Lewis bases are essential equivalent to Brønsted–Lowry bases. The only difference would be bonding to H+, versus bonding to some other Lewis acid. Bases have a pH greater than 7. Strong bases (and some concentrated weak bases) are, much like the acids, usually dangerous and corrosive. Remember, water auto-ionizes: H2O(l) + H2O(l) ⇄ H3O+(aq) + OH− (aq), so that Kw = [H3O+][OH−] This is really just KC for the above reaction, but since this reaction is important for water solutions, it’s given the special Kw symbol. At 25oC, the value of Kw is about 1.0 x 10−14. In pure water at 25 ⁰C, [H3O+] = 1.0 x 10−7 M and [OH−] = 1.0 x 10−7 M. Since large negative exponents are inconvenient, we can use –log to simplify them. Basically, p = -log. For example, pH = - log [H+] (or, pH = - log[H3O+]) The same is true for pOH and pKw. Since p = -log, pOH = -log [OH−] and pKw = -log Kw = -log (1.0 x 10−14). Combining all of these, Kw = [H3O+][OH−] = 1.0 x 10−14 Taking –log of everything, it ends up as: pKw = pH + pOH = 14. (at 25oC) The concentration of H+ and the concentration of OH− are tied together because of water’s auto-dissociation equilibrium. Here are some example values for [H+] and pH, and next to them, what [OH—] and pOH will be in the same solution. [H+] (molarity) 1 x 101 = 10 M 1 x 100 = 1 1 x 10—1 = 0.1 1 x 10—2 = 0.01 1 x 10—3 = 0.001 1 x 10—4 = 0.0001 1 x 10—5 = 0.00001 1 x 10—6 = 0.000001 1 x 10—7 = 0.0000001 1 x 10—8 = 0.00000001 1 x 10—9 = 0.000000001 1 x 10—10 = 0.0000000001 1 x 10—11 = 0.00000000001 1 x 10—12 = 0.000000000001 1 x 10—13 = 0.0000000000001 1 x 10—14 = 0.00000000000001 1 x 10—15 = 0.000000000000001 pH -1 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 [OH—] (molarity) 1 x 10—15 M 1 x 10—14 1 x 10—13 1 x 10—12 1 x 10—11 1 x 10—10 1 x 10—9 1 x 10—8 1 x 10—7 1 x 10—6 1 x 10—5 1 x 10—4 1 x 10—3 1 x 10—2 1 x 10—1 1 x 100 1 x 101 pOH 15 14 13 12 11 10 9 8 7 6 5 4 3 2 1 0 -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, whose combined colors that can be used to estimate the pH in the range 1-14. In this lab, we will use either universal indicator solution, pH paper, or a red cabbage solution to estimate pH. To get 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 that is exactly pH = 7. More ions in your solution makes a higher voltage difference. Also, the probe uses a special glass barrier that only responds to H+, so that it can 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. Procedure: Part I. Determination of pH of household substances. 1. Obtain the universal indicator and follow the instructions on how to use it. 2. Your instructor will prepare a series of test tubes containing solutions of pH 1-14 and will add some of the universal indicator to each. This will be your set of standards for the pH determination. 3. Obtain some household substances. If the substance is a liquid, place 1-2 mL in a test tube. Predict the pH of the solution before you measure it. Determine the pH with universal indicator paper (see note below). If the substance is a solid, grind some up, add a little water and determine the pH with indicator paper. After you have determined the pH with indicator paper, determine the pH with the red cabbage extract. Note: The correct technique for determining the pH with pH paper is to dip a clean dry stirring rod in the solution. Remove the rod and touch the rod to a dry part of the indicator paper. Compare the color of the paper to the color chart on the side of the indicator paper vial. Do not drop the indicator paper into the solution, as it will dissolve the indicators off of the paper. Part II. pH of a series of hydrochloric acid solutions. 1. Obtain 10.0 mL of 0.10 M hydrochloric acid and place in a clean dry 50.0 mL beaker. Predict the value of the pH. Measure the pH with the pH meter. Record the value. 2. Take 1.00 mL of the 0.10 M HCl(aq) in the previous step, and put it in another clean beaker. Add 9.00 mL of deionized water and stir. What is the new concentration of HCl(aq)? Predict the pH. Measure the pH. Record the value. 3. Add 90.0 mL of deionized water to the diluted HCl in step 2. What is the new concentration of HCl(aq)? Predict the pH. Measure the pH. Record the value. Part III. pH of acetic acid solutions. Remember, acetic acid is a weak acid, with a Ka = 1.8 x 10−5. 1. Obtain 10.0 mL of 0.10 M acetic acid and place in a clean dry 50.0 mL beaker. Predict the value of the pH. Measure the pH with the pH meter. Record the value. 4. Take 1.00 mL of the 0.10 M CH3COOH(aq) in the previous step, and put it in another clean beaker. Add 9.00 mL of deionized water and stir. What is the new concentration of CH3COOH(aq)? Predict the pH. Measure the pH. Record the value. 5. Add 90.0 mL of deionized water to the diluted CH3COOH in step 2. What is the new concentration of CH3COOH(aq)? Predict the pH. Measure the pH. Record the value. Data Sheet Part I Name:__________________________ Measurement of pH from the universal indicator. pH Color Observed pH Color Observed 8 1 9 2 10 3 11 4 12 5 13 6 14 7 Determination of pH of Household Substances Substance Predicted pH pH with indicator paper pH with universal indicator Data Sheet Part II Name:__________________________ pH of hydrochloric acid solutions Concentration of HCl(aq) Predicted pH pH with pH meter Predicted pH pH with pH meter 0.10 M pH of acetic acid solutions Concentration of CH3COOH 0.10 M Post Lab questions. 1. Calculate the predicted pH of the initial HCl(aq) solution. Compare this to the pH using the pH meter and calculate the percent error. 2. Calculate the predicted pH of the initial CH3COOH (aq) solution. Compare this to the pH using the pH meter and calculate the percent error. 3. How did the pH of the weak and strong acid solutions differ? Did the pH numbers for the weak and strong acids behave the same when diluted? If not, explain why. 4. How did the predicted pH of the household chemicals compare with the estimated values using the indicators? What did you discover about the chemicals around you?