Introduction
The goal of this lab is to apply principles of the ideal gas law to solve for n or the number of moles carbon dioxide produced, and compare the amount found using the ideal gas law to the actual amount. In order to complete this lab it's necessary to understand the apparatus below. By filling the Erlenmeyer flask completely full with water the mass of CO2 gas in the top of the flask can be determined. Since the combination of sodium bicarbonate and oxalic acid produces CO2 gas, this gas then moves from the gas generation bottle into the tube connected to the pneumatic trough. This gas then moves through the hole in the trough into the Erlenmeyer flask. The gas rises to the top and then pushes the water down and out. By measuring
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Calculate the average gas volume produced.
(325 mL + 219 mL) = 272 mL 2
3. Use Dalton’s Law of Partial Pressures to calculate the pressure with respect to only the carbon dioxide. Record the carbon dioxide pressure.
749 mmHg - 22.1 mmHg = 726.9 mmHg = 727 mmHg 4. Use the ideal gas equation, to solve for the number of moles of carbon dioxide produced. Remember that variables in this equation must have units of atm, L, and K. Also R = 0.0821 Latm/molK. Also, since you are interested in moles of CO2, use the CO2 pressure. Show your work for any conversions made in addition to your working for solving for moles. Record the number of moles of CO2. PV=nRT (0.956579 atm)(0.272 L) = n (0.0821)(295.8 K)
P = 727 mmHg ÷ 760 mmHg = 0.956579 atm (0.260189) = n (24.2852)
V= 272 mL ÷ 1000 mL = 0.272 L n = 0.0107 mol
T= 22.8oC + 273 = 295.8 K
5. Calculate and record the mass in grams of carbon dioxide produced. Convert moles to grams.
0.0107 mol = g ÷ 12.01 + 2(16.00)
0.0107 mol = g ÷ 44.01
0.0107 mol ÷ 44.01 = g g = 0.471
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By dividing the 0.471g (experimental value) by 0.599 g (actual value) and multiplying by 100 the percent recovered was found to be 78.6%. Since the percent recovered was less than the actual value the experiment lost some of the CO2 in the process. This means that the combination of human and experimental errors affected the results. One of the major experimental errors is knowing when the reaction between the sodium bicarbonate and oxalic acid is completed. By putting the stopper in the Erlenmeyer flask too soon not all of the gas will have made it into the flask. Additionally since we were using tap water the density is slightly off that of pure water. This slightly alters CO2 ability to take up space in the flask. The combination of these factors accounts for the percent lost in the
If the amount of baking soda is increased, then the amount of carbon dioxide produced will also increase up to a certain point, at which the amount of carbon dioxide will remain constant because the vinegar has become the new limiting reactant.
For these calculations assume the temperature remains constant at 25 °C. There are four steps here. First the initial concentration of CoCl2(iPrOH)2 needs to be determined. Second, the equilibrium concentration of CoCl2(iPrOH)2 must be found. Third is the calculation of the equilibrium concentration of CoCl2(MeOH)4, then fourth, the ratio [CoCl2(MeOH)4] / [CoCl2(iPrOH)2] at equilibrium. Show all work. Remember to state the equations you are using defining all variables and constants. Then substitute in the values and show the results. Be certain all sig figs and units are correct. You are expected to follow this format in all subsequent lab work without being prompted.
Evaluation I think that the method used in the experiment is not very accurate because the way we measure the amount of gas produced is not very
In my experiment, I will use an overall volume of 50 cm³ of 2moles of
The molar volume of the H2 in our experiment is very close to the theoretical molar volume, but I think that the deviation lies in the temperature of the H2O: in the first trial it is too high and in the second one too low.
In this experiment, there were several objectives. First, this lab was designed to determine the difference, if any, between the densities of Coke and Diet Coke. It was designed to evaluate the accuracy and precision of several lab equipment measurements. This lab was also designed to be an introduction to the LabQuest Data and the Logger Pro data analysis database. Random, systematic, and gross errors are errors made during experiments that can have significant effects to the results. Random errors do not really have a specific cause, but still causes a few of the measurements to either be a little high or a little low. Systematic errors occur when there are limitations or mistakes on lab equipment or lab procedures. These kinds of errors cause measurements to be either be always high or always low. The last kind of error is gross errors. Gross errors occur when machines or equipment fail completely. However, gross errors usually occur due to a personal mistake. For this experiment, the number of significant figures is very important and depends on the equipment being used. When using the volumetric pipette and burette, the measurements are rounded to the hundredth place while in a graduated cylinder, it is rounded to the tenth place.
The best way to measure gasses is by creating a closed system for an experiment
2. A student would make small waves, while another would time it with a stopwatch until it reaches the other side. This step would be repeated, however the wave would be bigger or smaller, in order to find out if changing the frequency, the speed would change.
There is also the potential of human error within this experiment for example finding the meniscus is important to get an accurate amount using the graduated pipettes and burettes. There is a possibility that at one point in the experiment a chemical was measured inaccurately affecting the results. To resolve this, the experiment should have been repeated three times.
I shall be measuring how much gas is given off. This will be done by measuring the amount of froth on the surface of the liquid. The oxygen released is collected in the form of these bubbles. The equation for the reaction is: (catalase) [IMAGE] H2O2 2H2O + O2 (hydrogen peroxide) (2 part water) (oxygen) I will change the concentration of H2O2 and O2 (making sure the volume stay the same, when one part of a H2O2 particle is taken, an O2 particle is added. Prediction
The purpose of the experiment is to identify and understand reactions under kinetic and thermodynamic control. A reaction under kinetic and thermodynamic control can form two different types of products. A reaction under kinetic control is known to be irreversible and the product is formed quickly. A reaction under thermodynamic control is known to require rigorous conditions. It is also reversible. The final product is more stable than the product made by kinetic control. The chart below shows the two types of reaction coordinates:
It makes Hydrogen Peroxide decompose into water and Oxygen. We will be measuring the amount of Oxygen released from the Hydrogen Peroxide. In order to do this we will use a measuring cylinder. This piece of apparatus measures the amount of Oxygen given off, measuring in cm3. This will help us measure the amount of Oxygen more accurately.
Carbon Dioxide is a colorless, odorless gas that occurs in small quantities in the earth's atmosphere naturally. The earth's ocean, soil, plants and animals release CO2. The formula of Carbon Dioxide is CO2. The CO2 molecule contains 2 oxygen atoms that each share 2 electrons with a carbon atom to form 2 carbon - oxygen double bonds. The atoms are arranged as so (OHT). This is called a 'linear molecule'.
Since the days of Aristotle, all substances have been classified into one of three physical states. A substance having a fixed volume and shape is a solid. A substance, which has a fixed volume but not a fixed shape, is a liquid; liquids assume the shape of their container but do not necessarily fill it. A substance having neither a fixed shape nor a fixed volume is a gas; gases assume both the shape and the volume of their container. The structures of gases, and their behavior, are simpler than the structures and behavior of the two condensed phases, the solids and the liquids
This will allow us to compare the actual value of CO2 produced and see whether the yeast metabolized the sugar aerobically or anaerobically. We had to use the given equations and the mole to mole ratios to see how much the theoretical value of CO2 was produced from 0.5 g of