chemistry post-lab

PS 141L: Drop the Base

Question of the Day {QOD}: How good are you at making a 3.00M HCl dilution given a ~6M stock solution?

Safety:

Splash

Goggles

COLLECT

HAZARDOUS

WASTE

Lab Apron

& Coat

Chemistry Concepts:

Acids, bases, indicators

pH

Equivalence Point

Mole Ratio/Stoichiometry

Techniques (refer to lab manual):

· General Safety

· Hazardous Materials

Lab Equipment and Techniques

Quantitative Skills

o Measuring Mass and Volume

o

Significant Figures

o

Titration

o

Data Analysis

o

Measuring pH

o

Data Manipulation

o

General Techniques

o

Gran Plot

Additional Resources: Check out the Chemistry Tutor DVDs at the Tutoring Center: Volume 3, Sections 5-13. Materials:

25mL or 50mL Buret

Storage bottle (labeled!!!)

 Ring stand with burette clamp

 25 mL and 5mL volumetric

pipettes w/ pump

100mL volumetric flask

Vernier LabQuest/pH meter

Stir plate

Sodium carbonate

Stir bar

6 M Hydrochloric Acid

 Balance & weigh boats

3 M Hydrochloric Acid

 Various beakers

Methyl orange solution

** SAVE YOUR 3M HCl Dilution … you will need it for next week! **

Discussion:

Acids and bases are some of the most important industrial chemicals used in the United States, accounting for approximately 265 billion pounds of materials produced or imported to the United States in 2012. They are used in a plethora of industries including pharmaceuticals, paper, flavors & fragrances, soaps, detergents, fertilizers, paints, and fibers.

The concept of acidic and basic (alkaline) materials stretches back as far as the ancient Greeks, Romans, and Egyptians. These societies identified and characterized materials according to properties such as taste and texture. Acidic materials have a sour taste and are commonly found in foods such as vinegar, citrus fruits, and sodas. Basic compounds often have a more bitter taste and are slippery to the touch. Bases are frequently employed in the production of soaps and detergents.

The first modern description of the chemistry of acids and bases came from Svante Arrhenius in 1884. He classified them in terms of their formulas and behavior in water where an acid was considered a substance that contains hydrogen and dissociates in water to hydrogen ions, H+, while a base was considered a substance that contains and produces the hydroxide ion, OH–, when it dissociates in water. This classification, while useful, did display some deficiencies.

First, this theory introduces the concept of a hydrogen ion, which is essentially just a proton, floating around in the very polar solvent, water. Because of the shape and electron density of water (a concept you will learn more about in lecture later this semester), there is a partial charge separation that makes the oxygen atom in water slightly negative and the hydrogen atoms in water slightly positive. Just as in magnetic interactions, opposite charges attract, so rather than free hydrogen ions floating in solution, a more reasonable theory shows this attraction resulting in the formation of the hydronium ion, H3O+. Therefore, a better way of looking at the Arrhenius theory is that an acid increases the concentration of hydronium when added to water while a base will decrease the hydronium concentration thereby increasing the hydroxide concentration.

Additionally, another key shortfall of the Arrhenius model is that it cannot accurately describe the chemistry of substances such as NH3 and K2CO3, which do not contain the hydroxide ion in their formulas but do produce a basic solution when dissolved in water. To overcome these shortfalls, Johannes Brønsted of Denmark and Thomas Lowry of England independently suggested new definitions of acids and bases in 1923 to remove these limitations. According to Brønsted-Lowry theory, an acid is a proton donor, any species that donates an H+ ion; an acid must, therefore, contain an ionizable hydrogen in its formula. A base under the new definition is a proton acceptor, any species that accepts an H+ ion; a base must contain a lone pair of electrons to form a covalent coordinate bond with the H+ ion. There are other models of acid-base chemistry, but for the purposes of this class a basic understanding of Brønsted-Lowry theory should be sufficient.

Today, acidity is typically determined by measuring the concentration of hydronium ions in solution and is frequently reported in units of pH. There are many different ways of reporting concentration, but for pH, you will use molarity which is defined as the moles of solute per liter of solution. Molarity therefore has units of mol/L and is abbreviated M. The pH scale, introduced in 1909 by Søren Sørensen, is a convenient way to express acidity based on the auto-ionization of water. The auto-ionization of water refers to the small but significant ability of water to react with itself in an acid-base reaction:

H2O (l) + H2O (l) ⇌H3O+ (aq) + OH– (aq)

In pure distilled water, the concentration of hydronium ions and hydroxide ions will be equivalent, each 1.0 x 10-7 M. The pH of a solution is defined as the negative base 10 logarithm of the concentration of hydronium, so the pH of pure water would be 7.0, which is defined as a neutral solution:

pH = –log10[H3O+] = –log(1.0 x 10-7 M) = 7.0

As the hydronium ion concentration increases, the pH value decreases and the solution becomes acidic. As the hydronium concentration decreases, the water becomes alkaline and the pH value increases. Typically the pH scale is represented as a continuum from 0–14. The image below shows the pH scale along with typical household and natural products and their corresponding pH values.

Because of the significant importance of acids and bases in chemical and biological processes, it is often necessary to know their concentrations as accurately as possible. In order to accurately determine the concentration of a compound, typically chemists will perform a volumetric analysis in which they use highly precise glassware such as volumetric flasks, volumetric pipets, and burets. The process used to determine the concentration of a solution with very high accuracy is called standardizing a solution. To standardize an unknown solution, you react that solution with another solution whose concentration is already known very accurately, typically using an experimental technique known as a titration.

mage result for titration indicatorWhen performing a titration, a known quantity (called an aliquot) of the unknown sample under investigation is typically added to a flask or beaker. This sample is called the analyte. Meanwhile, the sample of known concentration, called the titrant, is usually added to a buret. A buret is a very accurately graduated glass cylinder with a stopcock or pinchcock that allows the solution it contains to be delivered at any speed from a rapid stream to drop-by-drop. Note that if the sample of known concentration is a solid, it will usually be added to the beaker instead and the unknown would then be placed in the buret. As long as the appropriate data is collected during the experiment it is not terribly important which solution goes where. The titrant is added to the analyte until it reaches the stoichiometric point, the point at which the moles of acid equals the moles of base in the reaction. Titrations are typically monitored by either a pH meter or a chemical indicator. An indicator is a chemical added in a very small quantity to the analyte that changes color at different pH values. A table depicting a variety of acid-base indicators along with their expected color changes and pH ranges is shown below.

When using a pH meter, the equivalence point can be seen as the center of a very sharp slope signifying a drastic, sudden change in pH. When using an indicator, one should be chosen so that it changes color around the expected equivalence point of the reaction. The point at which the solution changes color in a titration is called the endpoint and, if the indicator was chosen appropriately, should correspond closely to the stoichiometric point (equivalence point). If an experimenter is relying solely on an indicator to determine the end of a titration, he or she must be careful not to “over titrate”. He or she should be looking for the point at which a single drop of titrant causes the analyte to change color. If extra titrant is added, the concentration of the analyte will not be calculated with a high degree of accuracy. Because the concentration of the unknown is dependent on how much titrant is delivered, it is pivotal to perform titrations with great care and precision to ensure accurate results. For a visual demonstration of how a titration is to be conducted, please watch this video.

Analyzing Results

Although properly performing a titration experiment is necessary for standardizing a solution, the experiment will ultimately be useless if the experimenter does not know how to analyze the results. Whether using a pH meter or a chemical indicator, the purpose of performing a titration is to determine how much titrant is necessary to completely neutralize the analyte. With that information, the starting concentration of the known, and the balanced chemical equation, it will be possible to determine the concentration of the unknown.

Example. A student needed to standardize a sample of sodium hydroxide for use in another experiment. To perform the standardization, she planned to neutralize it with a sample of potassium hydrogen phthalate, a weak solid acid with the chemical formula HKC8H4O4. For her first trial, she dissolves 0.453 g of HKC8H4O4 in 15.00 mL of distilled water and then she uses a pH meter to accurately record how the pH changes after each small addition of the unknown sodium hydroxide. She generates an unformatted titration curve that looks like the figure shown below:

mage result for weak acid titration curve

Note that in this experiment the curve starts at a low pH since the solution is acidic but then gradually increases as more base is added. As the reaction approaches the equivalence point there is a sharp jump in pH showing that the solution has switched from having excess acid to now having excess base. The middle of that jump is the actual equivalence point and for this experiment it has been determined to result from the addition of 24.1 mL of NaOH. If you were titrating an acid into a base, you should expect the curve to look like a mirrored version of this (starting at high pH and gradually decreasing until a sharp drop).

In order to calculate the concentration of NaOH from this experiment a balanced equation must be written:

NaOH + HKC8H4O4 NaKC8H4O4 + H2O

Since the mass of HKC8H4O4 was known exactly, the moles of acid can therefore be figured out using molar mass:

0.453 g HKC8H4O4

x

1 mol HKC8H4O4

=

2.22 x 10-3 mol HKC8H4O4

204.23 g HKC8H4O4

At the equivalence point, the stoichiometric moles of acid equals moles of base, so using the mole ratio from the balanced equation, the moles of base added can be determined:

2.22 x 10-3 mol HKC8H4O4

x

1 mol NaOH

=

2.22 x 10-3 mol NaOH

1mol HKC8H4O4

Since the moles of base have now been determined stoichiometrically and the volume of base has been found using the titration curve, the concentration of NaOH can now be calculated. Recall that molarity = moles/liter so the volume from the titration curve must be converted to liters (24.1 mL x 1L/1000 mL = 0.0241 L)

2.22 x 10-3 mol NaOH

=

9.21 x 10-2 M NaOH

0.0241 L NaOH

If the endpoint volume could not be easily determined from the experimental data, a Gran plot could be generated to fairly accurately find this value. The derivation of a Gran plot is a bit beyond the scope of this class, but essentially it is a graphical means of estimating the equivalence point volume of a titration using linear approximations of non-linear relationships, such as the one that exists between pH and titrant volume. An important caveat to the use of Gran plots is that these linear approximations are only valid near the endpoint. To generate a Gran plot for the titration of a base into an acid as in this example, the student should identify all of her data points 10% below her estimated equivalence point found from looking at the normal titration curve. In this case, she estimated the equivalence point to occur at 24.1 mL of base addition so she would take all of the data points from 21.69 mL of base added (90% of the perceived endpoint) up to 24.1 mL of base added. She would then take all of those volumes and multiply each of them by 10-pH where the pH in the exponent corresponds to the specific pH associated with each of those volumes. A plot of volume (on the x-axis) vs. volume x 10-pH (on the y-axis) should yield a linear graph. The x-intercept of that line tells you the Gran plot estimation of the equivalence point. If the student was titrating a strong acid into a weak base as you are doing in lab, the process would be the same except instead of plotting volume of base added vs. volume of base added x 10-pH, you would instead plot volume of acid added vs. volume of acid added x 10-pOH, where pOH = 14 – pH. A very good video illustrating how to generate a Gran plot in Excel can be found here.

For her second trial, she decides to look for the equivalence point by using an indicator instead of a pH meter. She begins by dissolving 0.471 g of HKC8H4O4 in 15.00 mL of water and adds 3 drops of phenolphthalein (an indicator that is colorless in acidic solution and magenta in basic solution). She does not refill the buret all the way in between trials so her buret reads 10.27 mL at the start of the experiment. She slowly titrates the base into the acid and finds that the solution turns a pale pink signaling the endpoint has been reached once the buret reads 34.19 mL. With just the starting mass of HKC8H4O4 and the two buret readings, she can calculate the concentration of NaOH for the second trial in the same fashion as for trial 1.

Moles HKC8H4O4 at the start of the reaction:

0.471 g HKC8H4O4

x

1 mol HKC8H4O4

=

2.31 x 10-3 mol HKC8H4O4

204.23 g HKC8H4O4

Moles NaOH required to neutralize all of the HKC8H4O4:

2.31 x 10-3 mol HKC8H4O4

x

1 mol NaOH

=

2.31 x 10-3 mol NaOH

1mol HKC8H4O4

Volume NaOH added:

Vadded = Vfinal – Vinitial = 34.19 mL – 10.27 mL = 23.92 mL = 0.02392 L

Concentration NaOH:

2.31 x 10-3 mol NaOH

=

9.66 x 10-2 M NaOH

0.02392 L NaOH

Therefore, the average concentration of the standardized NaOH was found to be 9.44 x 10-2 M.

If instead of a solid acid, the student used an acidic solution of known concentration, the calculations would be virtually identical, except she would calculate the moles of acid used by multiplying the volume used (in liters) by the concentration of the acid (in M). Once the moles of the known compound have been determined, the other calculations can be performed exactly as written. For more practice with titration calculations, please refer to this video.

Procedure:

1. Very carefully obtain about 50-60 mL of 6 M HCl stock solution in a small (100-150 mL) beaker

2. From this 6M HCl, transfer 50 mL to a 100-mL volumetric flask using a 25-mL volumetric pipet twice

3. Carefully and slowly dilute the HCl in the volumetric flask up to the fill line using a DI water bottle

4. Hold a gloved finger over the top of the volumetric flask and gently invert the flask a few times to ensure the acid is thoroughly mixed

5. Obtain a 25-mL or 50-mL buret

6. Rinse the buret with a small amount of tap water

7. Rinse the buret with a small amount of DI water three times (don’t fill the buret!)

8. Rinse the buret with about 5 mL of your dilute HCl rolling the solution around the buret to coat all sides and rinse out the leftover water droplets

9. Drain this rinse into an appropriate waste container

10. Fill your buret with the ~3M HCl

11. Drain a small amount through the tip to eliminate air pockets

12. Record the initial volume (note that it does not have to be 0 mL, it just cannot be above the 0 mL line)

13. Add ~0.8g sodium carbonate to a 250 mL beaker recording the exact mass added

14. Dissolve the sodium carbonate in ~25-30 mL of water

15. Add 3-5 drops of methyl orange to the beaker (the solution should be a yellow color)

16. Add a stir bar and place the beaker on a stir plate underneath the buret

17. Connect the pH Sensor to LaqQuest and choose New from the File menu

18. Use a utility clamp to suspend a pH Sensor on a ring stand. Position the pH Sensor in the sodium carbonate solution and adjust its position so it will not be struck by the stir bar. Turn on the magnetic stirrer, and adjust it to a medium stirring rate (with no splashing of solution).

19. On the Meter screen, tap Mode and change the data-collection mode to Events with Entry. Enter the Name (Volume) and Units (mL) and select OK.

20. Before adding HCl titrant, click and monitor pH for 5-10 seconds. Once the displayed pH reading has stabilized, click . In the edit box, type “0” (for 0 mL added). Click to store the first data pair for this experiment. You are now ready to begin the titration. This process goes faster if one person manipulates and reads the buret while another person operates the computer and enters volumes. WARNING: DO NOT HIT STOP UNTIL YOU TAKE YOUR FINAL MEASUREMENTS OR YOU WILL HAVE TO START THE TITRATION OVER AGAIN!

21. Slowly add the HCl titrant in 0.2 mL increments for the first ~3.5 mL, pressing and recording the volume on the LabQuest after each addition (this should be a running total of the amount of acid added, i.e. 0.00 mL, 0.20 mL, 0.40 mL, 0.60 mL, etc.)

22. Slow the addition down to 0.1 mL increments until the solution changes color, again pressing and recording the volume after each addition

23. Continue adding titrant in 0.2 mL increments for one mL past the endpoint, pressing and recording the volume after each addition

24. When you have finished collecting data, click . Pour the waste from the beaker in the provided hazardous waste container taking care to not lose the stir bar

25. Click on the graph tab and examine your titration curve to find the equivalence point volume

26. Export the data by either wirelessly emailing it to yourself or connecting it to the Lab PC and opening the data in the LoggerPro program. You will need this data to generate a properly formatted titration curve for your lab report

27. Titrate two more samples of sodium carbonate. If your first titration gave an appropriate curve using the LaqQuest, you may complete these titrations using only visual cues from the indicator. You may use the final volume in the buret from the previous trial as the initial volume in the buret for the next one (in other words, don’t refill the buret between trials unless it is almost empty) but be sure to record both the amount of acid you start the titration with and the amount you end the titration with so you can calculate the volume of acid actually dispensed

28. Once finished, dispose of the HCl remaining in the buret as hazardous waste and transfer the HCl from the volumetric flask into a properly labeled container for next week. ONLY DISPOSE OF THE USED HCl – YOU WILL NEED THE REST NEXT WEEK!

Post Lab: Experimental Results & Analysis (Cut & Paste Excel work into one document for submission)

Trial 1

Trial 2

Trial 3

Mass Na2CO3 (g)

Initial Buret Reading (mL)

------------------------

Final Buret Reading (mL)

------------------------

Volume HCl Added (mL)*

Moles Na2CO3 (mol)*

Moles HCl (mol)*

Concentration HCl (M)*

Average Concentration HCl (M)*

Standard Deviation Concentration HCl (M)*

Relative Standard Deviation (%RSD)*

Percent Error (Assume theoretical concentration is exactly 3.00 M)*

* A sample calculation must be shown for each row marked with an asterisk (*) to receive full credit. You do not need to include initial and final buret readings for the trial completed with the pH sensor. For that trial, write the equivalence point volume noted on the curve as your Volume HCl Added.

1. Complete the table above with all of your experimental data. (30 points)

2. Sketch the label you placed on your 3M dilution bottle and describe how you determined what information to put on the label. (5 points)

3. In Excel, organize the raw data from the standardization of your dilution using the pH Sensor. Attach this table to your post-lab report. Use proper formatting. (5 points)

4. Generate 1) a titration curve and 2) a Gran Plot of one of your trials. Use proper formatting. Label equivalence point(s) in both figures. Recall from the pre-lab reading that a good overview for how to do a Gran plot can be found here https://www.youtube.com/watch?v=xJ8UgdW-O8E however, you must use Va vs. Va x 10-pOH (recall that pOH = 14 – pH) (20 points)

5. Comment on the accuracy & precision of your experiment according to your calculations. If your results displayed high precision and/or accuracy, explain how you were able to achieve successful results. If your results displayed low precision and/or accuracy, explain how you could have improved your results. (10pts)

6. Prepare a statement that communicates your analysis of the results of experimentation related to the QOD. In other words, draw conclusions based on your data that answer the QOD (10pts). (Consider the following questions: Did you get the expected titration curve? Did you achieve the desired dilution concentration? How can you demonstrate statistically how well you performed the dilution?)

7. When conducting this experiment, the amount of water used to dissolve the sodium carbonate is never of concern; however the buret had to be rinsed with the experimental HCl solution prior to performing a titration for fear of excess water inside the buret. Explain why water added to the Na2CO3 has no effect on the data, whereas water added to the HCl solution may drastically affect the data (10pts).

8. When only using the indicator to track the progress of his experiment a student accidentally overshot the endpoint. How will this technique error affect the calculated concentration of the HCl? Be specific in your explanation (i.e. don’t just say more accurate, less accurate, or no effect)! (10pts)