Name_________________________ Team Name ______________________
CHM112 Lab – Acids and Bases – Grading Rubric Criteria
Points possible Lab Performance
Printed lab handout and rubric was brought to lab
3
Initial concentrations and pHs completed before coming to lab (fill out the concentrations and calculations on Data Sheet part II)
2
Safety and proper waste disposal procedures observed
1
Followed procedure correctly without depending too much on instructor or lab partner
2
Work space and glassware was cleaned up
1 Lab Report
All observations recorded
6
Question 1
1
Question 2
1
Question 3
.5
Question 4
.5
Question 5
2
Total
20
Subject to additional penalties at the discretion of the instructor.
Points earned
Acids, Bases and pH Theory: Acids and bases are very common in chemistry and even biology. We need to be able to identify acids and bases and understand their properties. Acids are named from the Latin “acidus” meaning sour. Acid solutions have a pH less 7, turn litmus paper blue and react with acids. 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 inorganic acids that we learned about in chapter 4 are Arrhenius acids. We can recognize them by the H in front of an anion with (aq). For example: HCl(aq) is an Arrhenius acid. The Brønsted-Lowry definition expands on this and defines an acid to be 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 compounds that we may not think of as acids like water and alcohols. The Lewis definition expands this further: a Lewis acid is a substance that can accept a pair of electrons (or a lone pair) to form a covalent bond. H+ does this when it reacts with OH– to make water. The H+ accepts an electron pair from the base to bond with it. Examples of Lewis acids include all Arrhenius and Bronsted acids as well as some metal cations. An Arrhenius base produces OH– in water. 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. CO has a lone pair and can act as a Lewis base. Please remember that concentrated strong acids and strong bases are corrosive and that some of the weak acids and bases are also hazardous! Always wear goggles when dealing with concentrated acids and bases. We know that water exists in equilibrium with hydronium and hydroxide (the autoionization of water.) H2O(l) + H2O(l) ⇄ H3O+(aq) + OH− (aq), with :
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. Kw will vary with temperature. In pure water at 25 ⁰C, [H3O+] = 1.0 x 10−7 M and [OH−] = 1.0 x 10−7 M. If the [H3O+] > [OH−], the solution is acidic and if [OH−] > [H3O+] then the solution is basic. We can also look at the pH. We use pH because 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 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. You must calculate the Molarity of the acid solutions as well as the theoretical pH before lab! 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 (what is your best guess?? Is the solution acidic or basic?)
pH with indicator paper
pH with universal indicator
Data Sheet Part II Name:__________________________ pH of hydrochloric acid solutions Please calculate the concentration of the diluted solutions as well as the pH of the original and diluted solutions before lab. Concentration of HCl(aq)
Calculated pH
pH with pH meter
0.10 M
pH of acetic acid solutions. Please calculate the concentration of the diluted solutions as well as the pH of the original and diluted solutions before lab.
Concentration of CH3COOH
Calculated pH
pH with pH meter
0.10 M
Post Lab questions. 1. Compare the calculated with the measured pH of the 0.10 M HCl(aq) solution. Calculate the percent error.
2. Compare the calculated with the measured pH of the 0.10 M CH3COOH (aq) solution. Calculate the percent error.
3. How did the pH of the weak and strong acid solutions differ? Did the pH values 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 pH values obtained using the indicators? What did you discover about the chemicals around you?
5. A student obtains a 10.0 mL of 0.100 M benzoic acid. The student dilutes the acid to a total volume of 100.0 mL. Calculate the pH of this solution.
