Bio Lab
Solutions & Dilutions, Acids & Bases – Breaking Bad Badge
(Adapted from Biology Laboratory Manual 10th Edition, by Darrell Vodopich & Hayden-McNeil Lab Simulations)
Solutions and Dilutions
Chemicals in living systems are in solution. A solution consists of a solute dissolved in a solvent. For example, salt water is a solution in which salt (the solute) is dissolved in water (the solvent). The concentration of solute in a solution can be expressed as either a percentage of the total solution as weight/volume OR as a measurement of the Molarity of the solution. We will work on figuring out solution concentrations using both of these methods as we will be using several different solutions over the course of the semester.
Percentage (weight/volume)
For the percentage method, the percentage of solute is the number of grams of solute per 100 mL of solution (weight/volume). It does not matter what chemical or liquid the solute or solvent are – this is strictly a percentage. For example, a 3% solution of sucrose is prepared by dissolving 30g of sucrose in 1L (1000 mL) of water (30 is 3% of 1000). If we wanted the same concentration of sucrose in final volume of only 100 mL, we would dissolve on 3g of sucrose in 100 mL of water (0.03 x 100 mL = 3g; 3g is 3% of 100). If we wanted 3% sucrose solution in a final volume of 500 mL; we would multiple 500 mL by 0.03 (3%), which equals 15. In order to make 500 mL of a 3% solution, we would dissolve 15g of sucrose in 500 mL water. Now – let’s practice:
1) How many grams of sugar would you add to 100 mL of water to make a 25% solution? ______________________
2) How many grams of calcium chloride would you add to 100 mL of water to make a 25% solution? ___________________________
3) How many grams of calcium chloride would you add to 500 mL of water to make a 25% solution? ___________________________
4) What percentage solution would you have if you mixed 5g of sugar in 500 mL of water? __________________________
5) If you only have 20g of sugar, what percent solution would you make if you dissolved it in 20 mL of water? ______________________ What percent solution would you have if you dissolved the 20g of sugar in 200 mL of water? ________ What percent solution would you have if you dissolved it in 2L of water? ____________________
Molarity
Molarity is the most common measure of concentration. It does matter what solute you use in this method of measuring concentration. The weight of 1 mole (M) of a chemical equals that chemical’s molecular weight in grams. A chemical’s molecular weight is the sum of the atomic weights of its component elements. For example, the molecular weight of water is 18g (2H = 2×1=2; O=16; 16+2=18). A mole of water weighs 18g. A solution with a molarity of 1.0M would be the molecular weight of the solute dissolved in 1L (1000 mL) of solvent. For example, a liter of solution containing 58.5g of NaCl (MW=58.5g) is a 1M solution of NaCl. Like in the percentage method, if your final volume is less than 1L, you would reduce or increase the number of grams of solute proportionally. For example, if you only wanted 500 mL of a 1M NaCl solution, you would dissolve half of the amount of solute that you would in 1L (58.5g/2=29.25g), since 500 mL is one half of the volume… Let’s practice!
1. How many grams of NaCl (MW=58.5g mole-1) would you dissolve in water to make a 2M NaCl solution in 1L final volume? _______________________
2. What is the molecular weight of Calcium chloride (CaCl2)? ______________ How many grams of CaCl2 would you dissolve in water to make a 1.5M CaCl2 solution in 1000 mL final volume? ______________________
In order to save space and time in the lab, scientist often make very concentrated solutions called stock solutions that they can later dilute with water to the molarity they need. This process is called dilution. The formula to determine how much of the stock solution is required for the desired molarity is:
Vi x Mi = Vf x Mf
Where Vi is the initial volume, Vf is the desired final volume, Mi is the initial molarity, and Mf is the desired final molarity.
For example, if you want to make a 1M solution of NaCl in 500 mL and your stock solution is 2M, then:
Vi x 2M = 500 mL x 1M
Vi = 500 mL x 1M / 2M = 250 mL
SO you would take 250 mL of the 2M NaCl and add 250 mL of water to get 1M NaCl solution in 500 mL.
Let’s practice determining concentrations and dilutions in the exercises below:
1. How many mL of a 6M NaCl stock solution should you add to make a 0.5M NaCl solution in a final volume of 250 mL? ___________________________
2. If you only have 150 mL of 2M NaCl stock solution, do you have enough solution to make 200 mL of a 1M NaCl solution? _________________________
3. How many mL of water would you need to make 250 mL of a 0.1M HCl solution? You have a 1M HCl stock solution already made. _________________________
Acids and Bases
There are millions of chemical substances in the world. Some of them have acidic properties, others, basic properties. Acids are substances which release hydrogen ions (H+) when they are mixed with water. Bases are substances which release hydroxide ions (OH-) when they are mixed with water. (This freeing of ions is called dissociation in both cases). Free hydroxide ions react with the hydrogen ions producing water molecules: H+ + OH- = H2O. In this way, bases diminish the concentration of hydrogen ions. A solution rich in hydrogen ions is acidic, a solution poor in hydrogen ions is basic, or alkaline. Some acids dissociate only in part and they are called weak acids; others dissociate completely, freeing large amounts of hydrogen ions. These are called strong acids. In the same way, the bases can be stronger or weaker. Diluted acids and bases are less concentrated and less aggressive in their actions. The acidic or basic degree of substances is measured in pH units. The scale used spans from 0-14. Substances with pH lower than 7 are considered acids, those with pH equal to 7 are considered neutral, and those with pH higher than 7 are considered bases. Substances with low pH are very acidic, while those with high pH are highly basic. Concentrated acidic and basic substances are very corrosive and dangerous.
pH is the measure of the concentration of hydrogen ions in a solution. As this concentration can extend over several orders of magnitude, it is convenient to express it by means of logarithms of base ten. As this concentration is always less than one, its logarithm always has the minus sign. To avoid having to always write the minus sign, it has been agreed to write this value with the positive sign. (This is the same as using the logarithm of the reciprocal of the hydrogen ion concentration). So, the pH is the logarithm of the concentration of hydrogen ions, with the sign changed: pH = -log [H+]. Thus, when pH has low values, the concentration of hydrogen ions is high.
There are substances which have the property of change their color when they come in contact with an acidic or basic environment. These substances are called pH indicators. Usually, they are used as dissolved substances, for example phenolphthalein and bromothymol blue. Often, to measure the pH special papers which have been soaked with indicators are used. These papers change color when they are immersed in acidic or basic liquids. This is the case of the well-known litmus paper.
Procedure 1: Determine the pH of Various Substances
1. Find 10 samples around your house to test. They need to be in liquid format in order to be tested. Some examples of what you can test are coca-cola, mouthwash, soap (make sure to dilute with water first!), milk, water, lemon juice, etc. Determine the pH using the commercial pH indicator paper provided to you in your at-home kit. Record your results in the table provided.
a. pH papers – Immerse an end of the paper in the liquid you wish to examine and remove it immediately. The pH of the liquid is determined by comparing the color of the paper to the scale of colors printed on its packet. Record the pH reading in the table below.
| Samples | pH by pH paper |
1. Which substance was the most acidic? Which substance was the most basic?
2. Were there any solutions that had a neutral pH? If so, which ones were they?
3. Were you surprised by any of your results? Why or why not?
Log-in to the Hayden-McNeil lab simulation website (http://courses.haydenmcneil.com) and click on the “Acids, Bases, and pH Buffers” simulation. Read through the background material provided and then click on the gray arrow at the bottom of the page. Open up the simulation by clicking on the green button. The simulation directions are available on the website and below. Use the simulation for Procedures 2-4.
Procedure 2: Introduction to pH Indicators
1. Take six small test tubes from the Containers shelf and place them onto the workbench.
2. Label the test tubes 1 – 6 by clicking on the tube and typing in the name and add a reagent to each test tube according to the table below. All reagents can be found on the Materials shelf.
3. Take a pH meter from the Instruments shelf and place it in the first test tube. To properly attach the meter, remember to place your cursor over the test tube when dropping the pH meter onto it. Repeat this procedure for test tubes 2 – 6. Record the color and pH of all six solutions to reference later.
4. Take a 50 mL beaker from the Containers shelf and place it onto the workbench.
5. Add 5 mL of bromothymol blue to the beaker.
6. Take a dropper from the container shelf and place it on the workbench.
7. Place the dropper into the beaker of bromothymol blue. You should observe the dropper filling with the liquid.
8. Using the dropper, add 2 drops of bromothymol blue to each test tube. Record the color and pH of all six solutions again after adding this material.
| Test Tube Number | Solution | pH before addition of bromothymol blue | Color before addition of bromothymol blue | pH after addition of bromothymol blue | Color before addition of bromothymol blue |
| 1 | 10 mL water | ||||
| 2 | 10 mL acetone | ||||
| 3 | 10 mL 5 M citric acid | ||||
| 4 | 10 mL 5% vinegar | ||||
| 5 | 10 mL 4 M ammonia | ||||
| 6 | 10 mL diluted bleach |
9. Clear your workbench by dragging instruments back to the Instruments shelf and by emptying containers in the waste bin and then placing the empty containers in the sink.
Why is bromothymol blue considered a pH indicator?
How accurate of an indicator is the bromothymol blue?
Does it change to a different color for every pH examined?
Does the addition of indicator change the pH of the solutions at all? Explain why or why not.
Based on your results with bomothymol blue, what color would a solution that was pH 13.5 be?
Procedure 3: Phosphate Buffer System
Part 1: Set-up
1. Take four small test tubes from the Containers shelf and place them onto one side of the workbench.
2. Label the test tubes 1 – 4 and fill the test tubes with solutions from the Materials shelf, according to the table below.
Note: The combination of the two phosphate solutions in the fourth test tube is the standard phosphate buffer.
Read the labels of the solutions on the shelf carefully, so you do not confuse the two phosphate solutions. Sodium hydrogen phosphate (Na2HPO4) and sodium dihydrogen phosphate (NaH2PO4) are not the same thing.
| Test Tube Set-up | |||
| Test Tube Number | Water | 0.1 M Sodium Dihydrogen Phosphate | 0.1 M Sodium Hydrogen Phosphate |
| 1 | 5 mL | ||
| 2 | 5 mL | ||
| 3 | 5 mL | ||
| 4 | 2.5 mL | 2.5 mL |
Part 2: Response to Hydrochloric Acid
1. Take a pH meter from the Instruments shelf and place it into test tube 1. Repeat for test tubes 2 – 4.
2. Record the contents of each solution along with its pH in the results table below.
3. Take a 50 mL beaker from the Containers shelf and place it onto the workbench.
4. Add 10 mL of 0.5 M hydrochloric acid (HCl) to the beaker.
5. Take a dropper from the Containers shelf and place it onto the workbench.
6. Place the dropper into the beaker of HCl. You should observe the dropper filling with hydrochloric acid.
7. Move the dropper onto test tube 1 and add two drops.
8. Repeat step 7 with the other three test tubes.
9. Record the new pH of each test tube in the results table below. Indicate whether it was a positive or negative change.
10. Clear your workbench by dragging instruments back to the Instruments shelf and by emptying containers in the waste bin and then placing the empty containers in the sink.
Part 3: Response to Sodium Hydroxide
1. Set up your workbench with four test tubes, as in Part 1.
2. Repeat the procedure outlined in Part 2, steps 1 – 10, using 0.5 M sodium hydroxide (NaOH) in your beaker instead of hydrochloric acid.
| Results from Procedure 3: Phosphate Buffer System | ||||||
| HCl Results | NaOH Results | |||||
| Test Tube Contents | pH | pH after HCl | Change in pH | pH | pH after NaOH | Change in pH |
| Water | ||||||
| NaH2PO4 | ||||||
| Na2HPO4 | ||||||
| NaH2PO4 and Na2HPO4 |
Questions:
1. Which solution was the least sensitive to the addition of acid or base?
2. Which of the four solutions tested is the best buffer against changes in pH caused by the addition of an acid or a base?
Procedure 4: Buffering Capacity of a Phosphate Buffer
Part 1: Addition of Acid
1. Take a small test tube from the Containers shelf and place it onto the workbench.
2. Add 2.5 mL of 0.1 M sodium hydrogen phosphate (Na2HPO4) and 2.5 mL of 0.1 M sodium dihydrogen phosphate (NaH2PO4) to the test tube.
3. Take a pH meter from the Instruments shelf and place it in the test tube. Record the pH of the phosphate buffer solution in the results table below.
4. Take a 50 mL beaker from the Containers shelf and place it onto the workbench.
5. Add 20 mL of 0.5 M hydrochloric acid (HCl) to the beaker.
6. Take a dropper from the Containers shelf and place it onto the workbench.
7. Place the dropper into the beaker of hydrochloric acid. You should observe the dropper filling with hydrochloric acid.
8. Add 1 drop of hydrochloric acid to the test tube.
9. Record the pH every time you add another drop of HCl in the table below.
10. Repeat steps 8 and 9 until the pH falls below 3 or until you have dispensed a total of 15 drops, whichever is reached first. Make sure that you record the pH after each drop is added.
11. Clear your workbench by dragging instruments back to the Instruments shelf and by emptying containers in the waste bin and then placing the empty containers in the sink.
Part 2: Addition of Base
1. Repeat the procedure outlined in Part 1, steps 1 – 10, using 0.5 M sodium hydroxide (NaOH) in your beaker. Add the base until the pH rises above 12 or until you have dispensed a total of 15 drops, whichever is reached first. Record results in the table below.
2. Clear your workbench by dragging instruments back to the Instruments shelf and by emptying containers in the waste bin and then placing the empty containers in the sink.
| Results from Procedure 4: Buffering Capacity of a Phosphate Buffer | |||
| Drops Acid | pH | Drops Base | pH |
| 1 | 1 | ||
| 2 | 2 | ||
| 3 | 3 | ||
| 4 | 4 | ||
| 5 | 5 | ||
| 6 | 6 | ||
| 7 | 7 | ||
| 8 | 8 | ||
| 9 | 9 | ||
| 10 | 10 | ||
| 11 | 11 | ||
| 12 | 12 | ||
| 13 | 13 | ||
| 14 | 14 | ||
| 15 | 15 |
Questions:
1. The strong acid and strong base used in this lab are the same concentration. Thus, the magnitude of the pH change caused by addition of a single drop will be the same on either side of the starting pH. Addition of acid lessens the pH and addition of a bases raises the pH.
Construct a graph of the pH versus the number of drops of acid or base added to the phosphate buffer. To show the relative magnitudes, use a negative value for each drop of acid and a positive value for each drop of base. For example, use -7 to write 7 drops of acid and 5 to show 5 drops of base. Copy the graph you construct and paste into this document to turn in.
2. Suppose you had a buffer containing 0.5 moles of sodium dihydrogen phosphate and 0.5 moles of sodium hydrogen phosphate. How many moles of hydrochloric acid would this phosphate buffer be able to accept before the pH of the solution began to change drastically?
3. Which chemical provides the conjugate base in the buffer containing NaH2PO4 and Na2HPO4?
4. Explain why phosphate buffer needs both NaH2PO4 and Na2HPO4 in order to resist changes in pH.
5. Predict what might happen if you made up phosphate buffer with only half as much NaH2PO4 compared to Na2HPO4.
6. Your lab mate attempts to use bromothymol blue to differentiate between two solutions – one that should be pH 7.3, and another that should be pH 6.7. What is your advice to your lab mate? Do you agree with his decision?
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