Titration Lab

Post a picture of your titration setup, and a photo of your analyte at the equivalence point. Write a short summary of your procedure. Calculate the percent ionization of vinegar and describe why it is such a low number.

Setup:

Following the directions on the lab sheet, we filled up the burette to the zero point with the 0.25 M NaOH base. We added vinegar to the flask along with distilled water and four drops of phenolphthalein. After securing the burette to the stand and placing the flask on the hot place, we started the magnetic stir of the mixture. Slowly, we opened the valve of the burette to release the NaOH base. Once the drops started to make more noticeably pink streaks in the mixture and it took longer to return to the clear color, we knew the equivalence point was approaching. At that point, we switched to a drop by drop method so that the mixture would not pass the equivalence point and become bright pink. After the drop that caused the mixture to turn light pink, we took the tip of the burette out of the opening of the flask and took the photos below. We also repeated this procedure for a second trial, knowing the relative amount of NaOH base that would be needed.

Trial one:

Ms. Kapinos looking pleased with our titration success:

Trial two:

Trial one:

Volume of NaOH used = 25.0 mL
pH of vinegar = 2.4
Molarity of acid = 0.86 M
Volume of acid = 0.00725 L
Molarity of base = 0.25 M
Volume of base = 0.0250 L

Trial two:

Volume of NaOH used = 25.5 mL
pH of vinegar = 2.4
Molarity of acid = 0.98 M
Volume of acid = 0.00720 L
Molarity of base = 0.25 M
Volume of base = 0.0255 L


Average concentration of vinegar = 0.875 M

[H3O+] = 10^-2.4 = 0.0040

Percent ionization = 0.46%

The percent ionization is such a low number because it does not have the properties that would allow it to break into ions. Strong bases have a 100% ionization because the acid and base completely ionize. Since vinegar is a weak acid, it does not completely break apart into ions.

pic creds to Ms. Kapinos #candid:

Solubility: A Guided Inquiry Lab

Introduction:

The purpose of this lab is to identify a solid by making use of its solubility curve. Solubility is the maximum amount of solute that with be dissolved in a specific solvent. Any solution under the curve is considered unsaturated, as more solute can be added. Any solution on the curve (or above) is considered saturated, as more solute can not be dissolved. My lab partner, Alison, and I made our own procedure that involved heating the solution of the salt and water in a bath and then stirring the solution to see whether the salt dissolved or not.

Procedure:

1) First, we measured the 5g of the salt using a weigh boat and scale.

2) Next, we filled a quarter of a beaker with faucet water. We also measured 10mL of distilled water in a graduated cylinder and then transferred that into a smaller beaker.

3) We added the salt into the smaller beaker with the 10mL of distilled water. Then we put the larger beaker on the hot plate to heat it up, and placed the smaller beaker into the larger one. We held the smaller beaker in place with a pair of tongs to be make sure no extra water from the larger beaker entered.

4) We placed a thermometer into the solution with the salt, and waited until the temperature reached 44 degrees Celsius.

5) Once the solution reached 44 degrees, we removed the solution to stir it. 

6) Since the solute dissolved, we repeated the process, this time adding 4g into the solution, so it had a total of 9g of salt and the same 10mL of water.

7) Since the solute dissolved again, we knew which salt it had to be, and cleaned up.

Data:

Trial one:

Mass of solute = 5g

Mass of solvent = 10g

Temperature of solution = 44 degrees Celsius

Trial two:

Mass of solute = 9g

Mass of solvent = 10g

Temperature of solution = 44 degrees Celsius

Solubility (from graph):

NaCl at 44 degrees Celsius = 3.7g
KNO3 at 44 degrees Celsius = 7g
NaNO3 at 44 degrees Celsius = 10.7g

After stirring the solution in the first trial, the solute dissolved completely. After stirring the solution in the second trial, the solute again dissolved completely.

Discussion/Conclusion:

The solute we had was sodium nitrate (NaNO3). Our claim can be justified when taking the solubility graph into consideration. If the salt was sodium chloride (NaCl), 5g of it in 44 degrees Celsius would cause the solution to be saturated, meaning not all the solute would have dissolved. If the salt was potassium nitrate (KNO3), 9g of it in 44 degrees Celsius would cause the solution to be saturated. However, after 9g of the salt were added, all the solute still dissolved. The only other salt it could have been was sodium nitrate, because its solubility curve was above those of the other two. 

As the temperature of a solution (with a solid as the solute) increases, the solubility of the solid also increases.

Alka Seltzer and the Ideal Gas Law

In this lab, we found the mass of carbon dioxide that filled up a balloon. When Alka Seltzer tablets react with water, carbon dioxide gas is produced. We filled up a test tube with water and put crushed Alka Seltzer tablets (powder form) into a balloon. Next, we stretched the opening of the balloon around the rim of the test tube and let the powder mix with the water, which caused the solution in the tube to fizz. Once the fizzing stopped (when the maximum amount of carbon dioxide was given off), we measured the circumference. Then, we found the volume of the space inside the balloon by filling it up to the same circumference with water, and measuring that amount. Using all these measurements, we were able to calculate the mass of the carbon dioxide.

Data:

Calculations:

Before powder and water mixed:

While powder and water mixed:

Jason lookin' $weg in the background:

Analysis Questions:

1) Discuss an area in this lab where experimental error may have occurred.

There were several areas for errors to have occurred in this experiment. For example, some of the crushed tablet powder fell outside of the balloon when we poured it in. Also, we may have measured the circumference of the inflated balloon inaccurately, as it was not a perfect sphere but rather a three dimensional oval shape. In addition, some of the water from the balloon fell out as we poured it into the graduated cylinder.

2) Choose one error from above and discuss if it would make "n" the number of moles of CO2 too big or too small.

If the balloon circumference that we measured is too low, the volume used in the ideal gas law equation is too low. As volume is directly related to the number of moles, this would cause the number of moles to also be too low.

3) Filling the balloon with water may be one place where error could have occurred. Using the value for the circumference of the balloon in cm, calculate the volume of the balloon mathematically.

r = 4.97 cm

V = 514.23 cm^3 = 514 mL

4) Compare your answer to #3 to the volume obtained by filling the balloon with water. Is it close? Which do you feel is more accurate and why?

No, the measured volume and the calculated volume are not close. The volume obtained by filling the balloon with water is 192 mL greater, and is more accurate. This is due to the fact that the balloon was not filled to a perfect sphere by the carbon dioxide gas (as seen above) nor the water. In both instances it was a three dimensional oval shape.

5) The ideal gas law technically applies to ideal gases. Give two differences between a real gas and an ideal gas. You may use your computer or book to research.

An ideal gas doesn't have attractions or repulsions among its particles whereas a real gas could have these interactive forces. An ideal gas also doesn't have a net loss of energy caused by its particles colliding, but a real gas could.

6) Would the CO2 you collected in this lab be considered ideal? Why or why not?

No, the carbon dioxide gas would not be considered ideal. This is because real gases aren't ideal, and ideal gases don't exist.

Advanced Questions:

1) Using the information provided on the label, and stoichiometry, calculate the mass of CO2 that should be collected. 

(We used two tablets instead of three)

2g citric acid ==> 1.374g CO2

3.832g sodium bicarbonate ==> 2.007g CO2

In this case, citric acid is limiting, so the theoretical yield of carbon dioxide is 1.37g CO2.

2) What percent is the percent yield for the CO2 collected in your sample?

1.26g/1.37g x 100 = 92.0%

3) CO2 is water soluble. The solubility around room temperature is around 90mL/100mL of water. What effect does that have on your calculated "n" value?

As carbon dioxide is water soluble, some of the gas could have dissolved in the water when we shook the test tube-and-balloon system. This would cause less carbon dioxide to fill the balloon, and the circumference and volume of the balloon would be too low. Therefore, our n value would also be lower.

Calories in Food Lab

In this lab, we figured out how much energy (in calories) was contained in three different foods: a cashew, pecan, and cheesepuff. We did so by placing each piece of food on a paper clip and lighting it on fire. On top of the burning food we placed a tin can, and on top of the can, a flask of water with a thermometer inside. Since the water reached its maximum temperature after the piece of food burned entirely, we recorded the thermometer reading after the flame went out. Using the mass of the water samples, the specific heat of water, and each change in temperature of the water to figure out the calories contained in the foods.

Before setting this poor cashew on fire:

While this cashew is on fire, burning slowly and peacefully:

Data:

 Burnt food:

 Alison and I:

Questions:

1) Did you measure a temperature change in the food sample or the water?

As there is no way to directly measure the temperature of a food sample, we measured the change in the water instead. This served the purpose of finding the change in temperature of the food.

2) Did you measure the energy released by the food sample or the energy gained by the water?

The energy released by the food sample is actually the same as the energy absorbed or gained by the water. By measuring the change in temperature of the the water, we were also measuring the energy released by the food. 

3) What happens to the small amount of energy that is not absorbed by the water?

The energy that was not absorbed by the water could've escaped through the holes in the tin can. In this expense of energy, there wasn't a completely effective transfer, as some of the heat released by the food was not absorbed by the water.

4) Were you surprised by any of the results? Explain.

I was not severely surprised by any of the results, but I found it interesting that the nuts contained more energy than the cheesepuff. It seems more intuitive to think that a cheesepuff has more Calories, or the "unhealthy stuff" as we know them to be, than healthy nuts.