Lab 14: Titration Lab by Katie, Maya, Alison and Meghana

In this lab we used a very stressful, difficult, and painstaking process called titration. We combined acid (vinegar) and a base (NaOH). An indicator, phenolphthalein, was used to tell when the solution went from its acidic form to basic form. At this point we would know the volume of the base and could use MV=MV to solve for the molarity of the acid. Then we used the pH of the base to find molarity of the base, and used this to find percent ionization. The ionization of vinegar was .456%. It is low because vinegar is a weak acid and so very few of its molecules ionize and become ions. 

We used a barrette to carefully combine the base and acid. The acid was distilled in water and we kept adding more base until the indicator showed that the solution had made the transformation. 

Lab 13: A Guided Inquiry Lab by Katie and Maya

Introduction:
In this lab we had to find the identity of an unknown salt. Salt have individual solubilities, so by dissolving them at different temperatures, we could tell what compound they were. We used water as the solvent, and dissolved the salt as the solute. We referred to the solubility curves to identify the solid.

Procedure:
First, we referred to the solubility curve of NaNO3, KNO3, and NaCl. We decided that if we dissolved 7 grams in 10 grams of 80 degree celsius water, then it cannot be NaCl because the graph shows that at 80 degrees 7 grams of NaCl would dissolve. After eliminating NaCl we had to test a point on the graph between KNO3 and NaNO3 so we added 7 grams to 10 grams of water at 30 degrees to test if it would dissolve. If it dissolves it then it is NaNO3 but if it cannot dissolve all 7 grams then it is KNO3.

Quantitative Data:
Trial 1- 10 grams of water
             7 grams of salt
             heated to 80 degrees celsius

Trial 2 - 10 grams of water
              7 grams of salt
              heated to 30 degrees celsius

Qualitative Data:
Trial 1- All of the salt dissolved so it is not NaCl

Trial 2- The salt did not dissolve therefore it is KNO3

Conclusion:
Because the 7 grams of salt did not dissolve in 10 grams of water at 30 degrees we knew that it could not be NaNO3 because it would have dissolved according to the solubility curve. Also, it is KNO3 because it 7 grams of salt dissolved in 10 grams of water at 80 degrees, which it would not have if it were NaCl according to the solubility curve. Therefore we narrowed it down to KNO3. For most solids, increasing the temperature increases the solubility. Keeping the temperature of the solution at constant was very challenging but we managed to keep it level during our experiment.

Heating the solutions in a hot bath!

Lab 12: Alka Seltzer and the Ideal Gas Law by Katie and Maya

We found the amount of CO2 produced by combining citric acid and baking soda. We did this by measuring the volume the gas took up, finding the pressure and using PV=nRT to find the amount of moles.

Data:
Mass of alka seltzer- 10.07 grams
Circumference of balloon- 38 cm
Volume of water that fits in balloon- 1066 mL
Room Temp. - 20 degrees celsius
Barometric pressure- 759.2 mmHg

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

The water residue that was left inside the balloon when we rinsed it may have mad our measurements inaccurate. Also we had difficulty transferring all of the alma shelter powder into the balloon so we may not have had all of the residue in the reaction.

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

Spilling the powder and not having all of it in the reaction may have caused our calculations to yield a lower the number of moles than what it actually was. 

3.  Calculate the volume of the balloon mathematically using the circumference you calculated in cm.

The actual volume of the volume of the gas that filled the balloon is 926.62

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

They were off but still relatively close. The calculation is more accurate because experimental measurements may have been inaccurate. 

5. List two differences between a real gas and an ideal gas.

Real gases have volume while ideal gases don't. Also, we assume that ideal gases have no intermolecular forces but in reality real gases do.

6. Would the CO2 you collected in this lab be considered ideal? 

No, because the gas that we found has volume, and although they are weak, they have intermolecular forces.

Advanced Questions:

1. The mass of the CO2 in grams was 2.062 g with the citric acid as the limiting reagent.

2. Our percent yield was 92.14%

3. Because some of the CO2 dissolved, the n-value we found may be a little less than the actual amount of moles. 

  

Lab 11B: Calories in Food by Katie and Zoe

We were able to find the number of food calories by solving for "chemistry calories". By burning the food and letting the heat be absorbed by water, we could solver the equation q=mcΔt where q lost by food is equal to q gained by water. Here is our data below:

Questions:

1. We measured the temperature of the water.

2. We measured both the energy released by the food and gained by the water because they are equal and set the two equations equal.

3. The energy is absorbed by the can or the beaker. It is absorbed into its surroundings.

4. I was surprised that the individual foods had such small amounts of calories, and that the cheese puff has less calories than the pecan and cashew. 

Lab 10: Evaporation and Intermolecular Attractions by Katie and Maya

2. Explain the differences in the difference in temperature of these substances as they evaporated. Explain your results in terms of intermolecular forces.

A larger differences in temperature means that more of the substance was able to evaporate. This means that the intermolecular forces are weak and readily break.

3. Explain the difference in evaporation of methanol, ethanol, and n-butanol. Explain your results in terms of intermolecular forces.

Methanol has the highest change in temperature, ethanol has the second highest, and n-butanol had the lowest. This means methane had the most evaporation with the weakest intermolecular forces, while ethanol had the less evaporations, and n-butanol had the least evaporation with the strongest intermolecular forces.   

4. Explain how the number of OH groups in the substances tested affects the ability of the tested compounds to evaporate. Explain your results in terms of intermolecular forces. 

The more OH in a compound, the more hydrogen bonds. This means that it is harder to break the bonds and have the substance evaporate which is why Glycerins temperature went up. 

Lab 7: Flame Test Lab by Katie, Roz and Maya

In this lab, we burned different compounds. The colors that they gave off were a result of excited electrons from the heat of the bunsen burner. Each element gives off a different color when burned so we could identify the the elements that the wooden sticks were soaked in.

PRE-LAB:

1. When electrons absorb energy in an atom, it will move to a higher energy level. This is the excited start. The grounded state is when the electrons are all at their lowest energy level.

2. Tp emit is to let off or discharge.

3. The excess energy that the atoms use comes from heat from the bunsen burner.

4. Different colors travel at different wavelengths. Each element emits different photons of different wavelengths. This is the emission spectrum.

5. We wash would wash the nichrome wires between each flame test so we don't burn a combination of elements.

Calcium

Copper 

Potassium

Unknown #1: Strontium Nitrate

Unknown #2: Potassium Chloride

We were able to identify these unknown substances by comparing the color to the colors of the known chemicals. When the colors of #1 burned red we knew it was Strontium Nitrate. When #2 burned purple we knew it was Potassium Chloride.

Electron Configuration: You're Never Too Old for Battleship

Don't be fooled by the board game. Electron configuration was a difficult concept to learn, but by the end of this game, we were pros. The most challenging part was when the lanthanides or actinides came into play and we had to count in the f orbitals. However this game helped me learn how to apply the noble gas shorthand when writing the configurations, and how to locate and element based on its electron configuration.

Happy one month anniversary to my camp mate and board game buddy even though all your
battleships were in the d-block

Lab 6: Mole-Mass Relationships Lab

In this lab, we created a chemical reaction by combining Sodium Hydrogen Carbonate and Hydrochloric acid. Then we separate the solids and hydrate, and try to find how much NaCl is left as a solid. 

Sodium Hydrogen Carbonate and Hydrochlorate Acid

We used a dropper to combine the acid and carbonate

We heated up the solution to separate the hydrates and solids

The NaCl that was left behind

 We extracted 90% of the NaCl that we should have. This was most likely because some was removed when we moved it using tongs, or when it was being heated, some of the NaCl popped out.

Lab 5B: Composition of a Copper Sulfate Hydrate Lab

We took a sample of carbon sulfate hydrate and tried to find the ratio of carbon sulfate and H2O.
We heated up the sample in an evaporation dish to find out how much water was in the sample and calculated the ratio.

Copper Sulfate Hydrate

Heating up the Copper Sulfate Hydrate

All the water has evaporated from the compound to create Copper Sulfate

1. Calculate the mass of the hydrate.

(Mass of evaporating dish and hydrate) - (mass of evaporating dish) = Mass of hydrate

45.88 - 44.97 = .91 grams of copper sulfate hydrate



2. Calculate mass of the water lost.

(Mass of evaporating dish and hydrate) - (Average of Mass of dish and hydrate after heating 2 times) = Mass of Anhydrous salt

(45.22+45.68)/2 = 45.45 = Average of Mass of the dish and hydrate after heating it twice

45.88 - 45.45 = .43 grams of water


3. Calculate percentage of water in the hydrate.

(Mass of Water)/(Mass of total hydrate)= percentage of water in the hydrate

(.43)/(.91) x 100 = 47% water 


4. Calculate percent error.

Actual percent is 36%               (|Experimental Value - Actual Value| / Acutal value) x 100 = % error

(|.47-.36| / .36) x 100 = 30.5 % error


5. a) moles of water evaporated

1 mole of H20 = 18.0148 grams
conversion:
(.43 g H2O/1) x (1 mol/18.0148 g H20) = .0238 moles of H20

b) Moles of CuSO4 that remain in evaporating dish:

1 mole of CuSO4 = 159.61 grams
conversion:
(.48 g CuSO4/1) x (1 mol/ 159.61 g CuSO4) = .003 moles

c) Find the ratio of moles of CuSO4 to moles of H2O.

.003/.003 = 1
.0238/.003 = 7.9 or 8
1 to 8

d) What is the empirical formula of the hydrate.

(CuSO4)1 (H20)8


However, because our percent error was fairly high, we predict a different result. Our percent of water lost is greater than the actual answer so the real coefficient for water should be smaller than eight.

Lab 5A: Mole Baggie Lab by Katie and Maya

In this lab, we received a plastic bag with an unknown substance in it. However we knew the mass of the empty plastic bag and the number of moles in the bag. From this information we had to find what substance was inside the bag.

We took the mass of the plastic baggie and and subtracted the mass of the empty plastic bag. WE found the mass of the substance in grams and divided it by the number of moles (0.045 moles) to find molar mass. Our five options were sodium chloride, potassium sulfate, zinc oxide, sodium sulfate, and calcium carbonate. After calculating the molar mass of these compounds, our result were closes to Sodium Sulfate.

Our bag A6 contained Sodium Sulfate.

Lab 4A: Double Replacement Reaction Lab by Katie and Maya

In this lab we experimented with different combinations of solutions to see if they form a solid precipitate. 

Even though Ms. K warned us about what the precipitate looks like I had expected a solid produced by a chemical reaction to be more structured and a little less fluid looking. However the precipitate caused the substances to become a little cloudier rather than clear. 

It difficult to be precise with the dropper! We were probably inconsistent with the amount of different chemicals in the reaction. This skill will probably come with time and practice. 

A precipitate can make the liquid look cloudy 

Here are all the equations of the chemical reactions with the net ionic equations listed below. 

At first balancing the chemical equations was also difficult but once I got used to the process it was easy to go through the steps of the process. 

Lab 3 Activity: Nomenclature Puzzle

The goal of this lab activity was to finish a puzzle matching the formula to the name of a chemical compound. For example one puzzle piece might say KNO3 and we had to match it to the puzzle piece that said Potassium Nitrate. This helped us practice naming compounds of Types 1, 2 and 3.

The biggest challenge was probably trying to locate all the matching pairs that were written in different directions. Also, there were so many! Once we had major pieces of the puzzle together, it became easier to locate these pairs.

My biggest contribution was matching together different formulas and names. Once we sorted the different elements, it became clearer which ones should be together, so I tried to match them in an efficient way.

Lab 2B: Atomic Mass of Candium by Katie and Nick

Purpose: The purpose of this lab was to try and find the average atomic mass of an "element" Candium. Candium has three "isotopes": regular, peanut, and pretzel m&ms. This helped us understand what the atomic number of an element really means and how to apply it when learning the different isotopes of an element.

Average Atomic Mass of Candium: 1.52 grams.

1. Ask a group nearby what their average atomic mass was.  Why would your average atomic mass be different than theirs?

We may have a different abundance of different isotopes. For example we had more m&ms with a larger mass than they did, so our average atomic mass went up. The abundance of each isotope affects the average atomic mass of an element.

2. If larger samples of candium were used, would the differences between your average atomic mass and others' average atomic masses be bigger or smaller? Defend your answer.

The differences would become smaller. As the sample increases, our abundances of each isotope is more likely to become closer together. Therefore our all of the average atomic masses become more accurate and closer to each other.

3. If you rook any piece of Candium from your sample and placed it on the balance, would it have the exact average atomic mss that you calculated?

No. The average atomic mass will not match the exact mass of any piece of candium. Since Average atomic mass is an average, it will be slightly higher or lower than any given isotope.

4. Here is the element square for candium!

Lab 2A: Chromatography Lab by Katie and Nick

Chromatograph #1 After the water

Chromatograph #1 before the water   

                                         

Chromatograph #2
Chromatograph #1

1. Why is it important that only the wick and not the filter paper circle being contact with the water in the cup?

The ink from the marker is being carried through the filter paper gradually by the water. This way we can see which substances are quickly absorbed into the paper and which ones can travel farther up the paper with the water. If the paper were fully submerged in the water, we would not be able to see the differences in distance travelled by different substances.

2. What are some of the variable that will affect the pattern of colors produced on the filter paper?

The distance from the center affected what kind of design appeared on the paper. We also learned that drawing various shapes and lines on the paper will change the design. The ink from different markers produced different colors. Somer were very colorful, others were just blue or black. Leaving the paper on the wick will allow the water to dilute the color more. For example, the bottom design was left in the water longer than the top design.

3. Why does each ink separate into different pigment bands?

Depending on what substance makes up the ink, some inks are easily absorbed into paper while others can flow through it easily like the water. From our observations, orange and yellow stay close to the middle because it stays printed on the paper. However, pink can travel farther than that and light blue is always out on the very edge.

4. Choose one color that is present in more than one type of ink. Is the pigment that gives this color always the same? Do any of the pens appear to contain a common pigment?

Light blue was a color used in many of the inks. Repeatedly light blue was the one that always spread out the furthest. The similarities between the chromatograms shows that different pens appear to use same pigments.

5.  Why are only water-soluble markers or pens used in this activity? How could the experiment be modified to separate the pigments in "permanent" markers?

Water soluble markers can be broken down when the water passes through the ink. Therefore, it is easier to see the separation of the pigments when they spread out on the adsorbent along with the water. Using a solvent that can break down the different pigments in the ink of a permanent marker will allow us to create a chromatograph using a permanent marker.

    

Lab 1B: Aluminum Foil Lab by Katie and Maya

Procedure: First we knew we could find the thickness of the tinfoil by using Density=mass/volume. We found the length and width of the tinfoil using a ruler and rounding to 3 significant figures (in cm). Then we found the mass of the tinfoil using a scale (mass found in grams), and found the density using the water displacement method.

The water displacement method is using a graduated cylinder to find the volume of an irregular object. The water was at first at 89.0 mL and after the tinfoil was submerged under the water, the water level rose to 90.0 mL. The difference in the water levels was 1 mL and so the volume of the tinfoil is 1 mL.


After finding the density, we used our knowledge of the relationship between volume, mass, and density to calculate the thickness of the tinfoil.

Lab 1A: Density Blog Lab by Katie and Maya

Introduction: Using what we learned in class about density, we could find the mass of a block using volume and density. Density, which is the degree of compactness of a substance, can be found by dividing mass (g) by volume (cm^3). Using this knowledge, our knowledge about significant figures, and a ruler, the mass of each block could be found without a scale.

Procedure: Using a ruler, we measured the length, width, and height of the blocks in centimeters to three significant figures. Multiplying the length, width and height together we could find the volume.

From there we could use the equation d=m/v and multiply volume and density to find the mass of the block which ended up being 94.5 grams. The actual mass of the block ended up being 97.3 grams so our percent error was 3%. 

Data: 

Length-6.70 cm

Width-4.19 cm

Height-  2.37 cm

Volume-66.5 cm^3

Density- 1.42 g/cm^3

Mass- 97.3 grams

Predicted Mass- 94.5 grams

Percent error- 3%

Conclusion: In this lab we learned the relationship of volume, density, and mass. We learned that you can find one from knowing the two others. We also used a ruler while keeping in mind our significant figures. There was, however, a margin of error when measuring the dimensions of the block. It was difficult to estimate the last significant figure for the measurements. I realized that next time, it may be better to combine my estimate and my lab partners in order to get a more accurate measurement.