Calorimetry Lab

Procedure: I used a ruler, thermometer, and scale to take measurements. I used a graduated cylinder, short step pipet, scale, and ruler to determine volume and density. I used a volumetric flask, graduated pipet, pipet bulb, scale, and glass beaker to determine concentrations and densities of various dilutions.

After the twenty minutes elapsed, the flask was cooled to room temperature and then titrated with the remaining NaOH until the colorless solution remained pink. The final volume was then recorded. While solution #1 was heating the same process was repeated with solution#2 and the second burette

When filling up the burette it is important that a funnel is used, however as the solution reaches the 0 mark it is ideal that the funnel be removed and a pipette used instead to reach the 0 mark, this is to achieve greater precision. During the experiment, it is important to swirl the flask continuously with one hand

Identifying an unknown substance can be accomplished with the use of multiple scientific tests, which help narrow down the possibilities of the unknown compound. The unknown substance that I was given was unknown number A84841BIR, and posed to be a real challenge since I needed to test two different molecular combinations for this specific substance. Once the number of moles in substance was calculated using the ideal gas law, I divided it from the sample mass number after heating and obtained that the molarity of my unknown chemical was 58.93 g/mol. To find my molecular formula I calculated the mass of each element, then the moles, lastly I took the smallest whole-number ratio. Using the molarity calculated I was able to do some research to get a better understanding about what my unknown chemical could possibly be. Once I’ve collected all my data together and strenuously researched the web for articles relating to understanding Infa-Red spectroscopy, I should be able to make an educated guess upon what my unknown solution could be.

made supported one another. For instance, we observed that the red powder turned black when heat was applied, therefore Oxygen from

The first two parts of the experiment looked at solubility, where part one looked at organic compounds and part two looked at alcohols. The solvents used in the first part were water, methanol, and hexane; decreasing in polarity respectively. Compounds that have the same polarity (both polar or both nonpolar) tend to dissolve when mixed with each other. This is shown when malonic acid (C3H4O4) is mixed with water (12 seconds) or methanol (instantly) and only one layer of liquid is formed. Also, malonic acid mixed with hexane formed two layers (insoluble) supporting the claim a nonpolar compound (hexane) and a polar compound do not readily dissolve.

Steps 1, 2,6,7,8,9,10 and 12 are all to do with measuring. This is an important factor when thinking about trying to get as much yield as possible because if the accuracy is wrong we would lose yield and therefore would reduce the final product. For example, during step one, the alcohol has to be poured into the pear shape flask. There is a chance that impurities could attach to the alcohol when being measured out, for instance dust particles going into the flask as well. This could therefore contaminate the yield and would reduce the purity. Also, when pouring the alcohol into the flask,

When the surfactant concentration exceeds the critical micelle concentration (cmc) in water, micelles are formed as aggregates of surfactant molecules. [6]

The mass of the distillate was obtained after fermentation and the fractional distillation was calculated to be 0.195 g. The distillate was collected in three fractions using a 1 mL syringe with a needle when the temperature reached approximately 78°C. The amount of ethanol that should have been produced was 1.08 g. However, this experiment only obtained 0.195 g of ethanol product. The percent yield was calculated to be only 18.12%, so only 18.12% of ethanol product was recovered from the fermentation reaction. The density of the distillate obtained from flask one was 1.03 g/mL, flask two was 0.979 g/mL and flask three was 0.99 g/mL. The distillate in all three flasks contained mostly water and only some ethanol because the percent recovery

The purpose of the fractional distillation of alcohols lab was to learn a new technique to separate mixtures of our unknown compound into pure components using the specific vapor pressures of pure liquids. Before the fractional distillation lab was performed, a simple distillation lab was run. In the simple distillation setup, a boiling flask attached to an adapter holding a temperature probe. The adapter connected to a condenser into which water passed through. The condenser leads to a flask for the purified liquid. Fractional distillation was a better choice for this lab over simple distillation because

The six boiling tubes were labelled A-F and then each was filled 10cm3 of starch solution with the help of a syringe.

Measuring the rate of osmotic pressure between tap water and surcose solutions. This will happen by using different concentrations of sucrose (20%, 40%, 60%) afterwards, the change in mass ratio would be observed over time.

Where B is the buoyancy force, Pb and Pt is the pressure at the top and bottom of the object respectively, A is the surface area of the object, Dfluid is the density of the fluid, g is the acceleration of gravity (9.8 m/s^2), h is the height of the object, V is the volume of the fluid displaced by the object, and M is the mass of the fluid displaced by the object.

In the first part of the experiment, the stock solution was prepared with a known initial volume, a known initial concentration, and a known final volume. The final concentration was calculated using the dilution equation, which showed how dilute or how concentrated the solution was as the experiment progressed. To form a solution at a certain concentration, such as a more diluted solution for an experiment, a solute is dissolved in a solvent thoroughly, and brought to the required volume of solution. Throughout the experiment, this was done repeatedly. In the first part of the experiment, as table 1 states, the final molarity concentration of test tube 1 was .10 M. As it progressed and the dilution equation was used to calculate the final concentrations of the new solutions, test tube 2 had a final molarity of .01, test tube 3 had .001 M, test tube 4 and 5 having .0001 M and .00001 M respectively. With each test tube, there was a progressively less concentrated initial molarity of the solution, which resulted in the less concentrated final molarities. For example, test tube 1 which began with an initial molarity of .10 M, and ended with a final molarity of .10 M, is significantly more concentrated than test tube 5, which had an initial molarity of .0001 and a final molarity of .00001. The results show that as the smaller molarity of initial solution was used, the smaller the final molarity was. This was also evident through the color change throughout the test tubes. The

Another source of error that was present in the investigation was rinsing the flasks between use. For example, the team members will rinse out the flask after testing it into the sink and then rinsing the insides with distilled water. However, it was difficult to dry the insides, so when reusing the flask, there may be extra residue of distilled water