Use the following as a checklist each time you finish a lab to ensure that you are submitting a complete lab report.
Introduction: (1 point each)
□ Background information – based on the concepts covered in the lab
□ Purpose – usually stated in the online write-up
□ Hypothesis and/or prediction – this one may not be present in all labs
□ Explanation – required if there is a hypothesis/prediction; otherwise, may be included in the background information
Procedure: (2 points each)
□ List of materials – be exact, particularly where important, such as “Fisher electronic analytical balance, +0.01 g,” not “scale,” or “two 12x150mm glass test tubes,” not simply “test tubes”
□ Numbered step-by-step procedure – paraphrase from the lab procedures (do not cut and paste)
Data – (points vary with lab)
□ Tables and/or Figures(graphs)
□ All tables/figures must be clearly labeled. The title should be specific enough to allow any reader to know what the figure/table is showing without reading the lab.
□ Proper labels must be included for the columns and rows or x- and y- axes. Here are two examples “Length/cm + 0.01” or “Temperature/oC + 0.1.” The intervals on the axes should be equivalent.
□ Numbers only go in the body of the table itself; for figures that are graphs, be sure that you use bar or line graphs as appropriate.
If you are instructed to download a data table, then you must make sure that it is titled and labeled properly. Use the graphical analysis program for figures that are graphs, and be sure to follow the above steps.
Discussion- (points vary with lab)
□ Answer the questions. These questions are called “Analysis,” “Calculations,” “Discussion,” or “Processing the Data” (depending on the lab). Make sure to number them according to the way they are numbered in the lab.
Introduction: Connecting Your Learning
In acid-base equilibria, some substances can act as either an acid or a base; that is, they can either increase the hydrogen ion [H+] or hydronium ion [H3 O+] concentrations, or increase the hydroxide ion concentration [OH-]. For example, amino acids exhibit this property as follows:
The general formula of an amino acid is H2N-(CHR)-COOH, where R represents the rest of the specific amino acid, such as another H for the amino acid glycine, which is H2N-(CH2)-COOH.
Amino acids are the basic building blocks of all proteins.
H2 N(CHR)COOH(aq) +H2 O(l) ←→ H2 N(CHR)COO- (aq) +H3 O+ (aq)
H2 N(CHR)COOH(aq) +H2 O(l) ←→ + H3 N(CHR)COOH(aq) +OH- (aq)
The equilibrium can be shifted by the pH; namely, by the [H+](aq) or [OH-](aq). Such substances can help prevent wide swings in the pH of a system, which can be critical for the biochemical reactions in living organisms. For example, human blood pH must remain between 7.35-7.45 or serious malfunctions can occur, including death. There is another way to reduce rapid changes in pH. For example, what is buffered aspirin? Why are buffers added to things ranging from shampoos to contact lens solution? A buffer resists a change in pH. Aspirin, which is acetylsalicylic acid, can irritate a stomach by increasing the [H+] (lowering the pH), otherwise known as acid indigestion. A buffer can react with the H+ to slow the lowering of the pH. Hair is made up of protein, which is made up of amino acids. The alkaline (the amine part) and acidic properties of those amino acids give the protein the shape it forms. Since shampoos tend to be alkaline, unwanted changes can occur chemically with the hair if the pH significantly changes. Buffering the shampoo (adding ingredients as a chemical buffer) can reduce those unwanted side effects. Finally, the surface of the eye is covered in part by proteinaceous mucus. Therefore, any chemicals you put in the eye should likewise not cause drastic pH changes. By including a buffer with the contact lens saline solution, this can be achieved, which is especially important since some salts can change the pH.
A buffer works by being able to neutralize both H+ as well as OH- . In the case of a buffer made with a weak acid, it can do this by both being able to provide H+ (to neutralize the OH- ) as well as being able to bond with the H+ (due to the fact that the H+ of the weak acid will tend to “stick” to the anion; that is, much of the weak acid does not ionize). Here is an example of a buffer system and how it would work:
Buffer system: CH3 COOH(aq) +CH3 COO- (aq) ←→ H+(aq) +2CH3 COO- (aq)
CH3 COOH(aq) (ethanoic acid or acetic acid-vinegar) is the weak acid and CH3 COO- (aq) (ethanoate ion or acetate ion) is the anion. Either the weak acid can dissociate into H+ and the anion (CH3 COO- ), which shifts the equilibrium to the right, or the weak acid can reform into the molecule (CH3 COOH) to shift the equilibrium to the left. In general, a buffer system is HA + A-, where HA is the weak acid and A- is the corresponding anion. (For a weak base, the system looks like this: B(aq) +H2O(l) +BH+(aq) ←→ 2BH+ (aq) +OH- (aq) , or in general, B + BH+ .)
The way the buffer neutralizes both H+ as well as OH- is as follows:
Adding acid (H+) to the above buffer system-
H+(aq) +CH3COOH(aq) +CH3COO- (aq) →2CH3COOH(aq)
Adding base (OH- )-
OH-(aq) +CH3COOH(aq) +CH3COO- (aq) →H2O(l) +2CH3 COO- (aq)
Adding more acid uses up the hydrogen ions by creating more of the molecular ethanoic acid, while adding more base uses up the hydroxide ions by creating water. Basically, the addition of H+ or OH- simply shifts the buffer system to the left or right. As long as there is sufficient anion to neutralize the acid as well as sufficient weak acid to neutralize the base, the buffer will slow the change in pH.
Buffers are made by adding a weak acid (or weak base) to a salt of the weak acid (or weak base). For example, adding sodium acetate (sodium ethanoate, NaCH3 COO) to ethanoic acid (acetic acid, CH3 COOH) will result in the following:
NaCH3 COO(aq) +CH3 COOH(aq) ←→ Na+ (aq) +CH3 COO- (aq) +CH3 COO- (aq) +H+(aq)
Since sodium acetate completely ionizes in water, you essentially have this:
Na+(aq) +CH3 COO- (aq) +CH3 COOH(aq) ←→ Na+(aq) +2CH3 COO- (aq) +H+(aq)
The Na+(aq) is a spectator ion since it is on both sides of the equilibrium and does not participate in any of the reactions. Therefore, what you really have is the exact buffer system that was listed as the first example earlier. The more weak acid you add, the more base can be neutralized, obviously, and the more salt of the weak acid you add, the more acid can be neutralized. In this lab, you will be creating and testing a buffer.
Resources and Assignments
Lab 6 Report
Materials (Lab Kit)
9 Dropper pipettes
4 Sample cups
Na2 CO3 (Washing soda) – premassed sample
NaHCO3 (Baking soda) – premassed sample
Standardizing buffer solutions
Materials (Student supplied)
Clear drinking glasses (12 oz.)
Focusing Your Learning
By the end of this lesson, you should be able to:
Determine the controls and experimental groups in a lab.
Apply the Brønsted-Lowry and net ionic equations concepts.
Evaluate a simple buffer system.
Determine an amphoteric system.
Print out a data table for this experiment.
Lay out newspaper or paper towels over your work area.
Create a 1.0 M solution of sodium carbonate using the premassed sample of this solid in your lab kit, and distilled water. Mix the ingredients together in a clean glass. Since you will have to use the entire premassed sample, what you will have to determine by calculation prior to mixing, is how much distilled water to add to the solid sample of sodium carbonate. (You may want to check your calculations with your instructor before beginning to ensure that you are creating the solution appropriately.) Measure the water accurately to the nearest tenth of a milliliter using the graduated cylinder. Stir until solid is completely dissolved and mixture is clear. This may take several minutes.
Repeat Steps 1 and 2, using NaHCO3 to make a second solution of 1.0 M sodium bicarbonate. Use a second clean glass to make this solution.
Label each glass and place a labeled dropper pipette into each solution, as these pipettes will be used to distribute the solutions during the experiment.
Label another glass and pipette for distilled water.
Label two more pipettes for the vinegar and ammonia (to avoid cross-contamination between solutions) and label two sample cups for these two substances. Fill each cup about half way with the indicated solutions.
Set up four (4) sample cups, and label them a, b, c, and d. The four (4) cups will each have different solutions for testing in them. (Rinse the graduated cylinder with distilled water in-between measuring each of the liquids.) Add to the cups the following chemicals:
Into cup a., add 10ml of distilled water
Into cup b., add 10m. of Na2 CO3 solution
Into cup c., add 10ml of NaHCO3 solution
Into cup d., add 5ml of Na2 CO3 solution + 5mL of NaHCO3 solution
**You will want to recalibrate your pH meter before performing the lab. Please follow the directions provided in Lab 5 for calibrating your pH meter.
Review the video on calibrating the HI98103 pH Meter.
Using the pH meter, determine the pH of all four (4) solutions. Rinse the electrode using tap water in between measuring the pH of each different solution to avoid contamination between solutions.
Do not throw out buffer solution!! You will want to fold over and seal the solution for Lab 7.
Using the vinegar dropper, add two (2) drops of vinegar to each cup, and swirl the cups to mix each solution.
Using the pH meter, determine the pH of each solution after the vinegar has been added and mixed in. Rinse the pH meter using tap water in between measuring the pH of each solution. Record both the total number of drops of vinegar added to each solution, and the pH, in the data table.
Repeats Steps 11 and 12 until 12 drops of vinegar have been added in two-drop increments to each of the four (4) solutions in the sample cups. Measure and record the pH of the solutions after every two (2) drops of vinegar has been added.
You can pour the contents of the cups down the drain, and rinse the cups first with tap water and then again with distilled water.
Set up the cups again, according to Step 8, using the four solutions. Test the initial pH and record the pH in the data table.
Add one (1) drop of ammonia to each cup (a, b, c, and d) and swirl the cup to mix each solution.
Using the pH meter, determine the pH of each solution. Rinse the electrode in tap water in between measuring the pH of each solution. Record the total number of drops of ammonia in solution and the pH in the data table after each measurement.
Repeats Steps 16 and 17 until six drops of ammonia have been added in one-drop increments. Record the pH of the solutions after every additional drop of ammonia is added.
Assessing Your Learning
Conclusion and Discussion Questions
Write the formulas for both the weak acid and conjugate base used in the buffer. Write the dissociation reactions for each substance. (See section 16.2 in your textbook on Brønsted-Lowry Acids and Bases.)
If each trial of this experiment was controlled, what were the controls?
Within each individual trial, (cups a, b, c, and d) which test tube served as a control group?
You know that the buffer system was the baking soda and washing soda together. Why did you test these substances individually (cups b and c)?
When you tested the baking soda and washing soda individually, did either of them work alone as a buffer? Explain why you would or would not use empirical (experimental) evidence in your response.
The acid you used was vinegar, and the base you used was ammonia. Which substance caused a greater change in the pH of each system, and why do you think this is so? (Keep in mind that the pH of pure vinegar is about three and pure ammonia is about 12).
Write the total ionic equations for a) the addition of the acid to the buffer, and b) the addition of base to the buffer.
Write the net ionic equations for a) and b) from the previous question.
Buffers are only useful over a certain pH range. For what pH is this buffer solution useful?
The buffer solution you made had nearly equal molar parts of weak acid and conjugate base. If you wanted to use this same buffer system (bicarbonate- carbonate) at a lower pH (more acidic), how would you need to change the proportions of weak acid and conjugate base to lower the pH of the buffer solution?
How would you change the proportions of the weak acid and conjugate base to make a buffer solution that is useful at a higher pH (more basic)?
What substance in this lab is amphoteric?
What empirical evidence do you have that this substance has amphoteric properties?
An unknown strong acid or strong base spill (you don’t know which) has occurred in the laboratory. You have at hand solid potassium hydroxide, sodium hydrogen carbonate, water, sand, and concentrated hydrochloric acid. Which of these should you sprinkle or pour onto the spill to neutralize it? Explain.
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