Data: Data Table 1 | | | Rubbing Alcohol Trial # | Boiling Point | Percent Error | 1 | 79.5 C | 3.52% | 2 | 84.0 C | 1.90% | 3 | 83.5 C | 1.30% | Data Table 2 | | | | Acetamide Trial # | Melting Point | Freezing Point | Percent Error (Melting Point) | 1 | 79.5 C | 57.5 C | 3.50% | 2 | 80.5 C | 59.5 C | 0.60% | 3 | 78.0 C | 61.0 C | 3.70% | Observations: It was observed that thermometer placement could affect the readings in the water bath. I ended up holding the test tube assembly to where the bottom of the test tube was not touching the bottom of the beaker to ensure better readings. Also, when substances recrystallize, I am not sure whether to take a reading when the substance has full crystallized or begun to so results could be fairly inaccurate. Questions: A. Why is it useful
The products of a combustion reaction are carbon dioxide and water. I also found out that it takes skill to light a match and that it’s difficult to keep gas where you want it to be without a sealed container. There are many sources of error. The first source of error was that if you blow on the candle too hard, the wax vapor blows away and the flame doesn’t jump. The next source of error is that if you don’t put in the limewater and stopper quickly enough, the carbon dioxide escapes and the water won’t turn chalky.
There was a low efficiency rate for this experiment; energy was most likely wasted into the surrounding environment when the burner was alight. Possible ways to improve this experiment would be to possibly do the experiment in a more enclosed space, so as to disallow any heat escaping into the surrounding atmosphere. A fume cupboard would be suitable (when it is not turned on) as there is less movement in the air to move the energy from its intended target. The thermal energy was not only going into the water, but the can of the calorimeter became hot too, meaning that the thermal energy was transferred into the metal surrounding the water, and not just the
Set the stopwatch to 1 minute. Use a lighter to light the fuel. Blow out the fire when the timer reaches 1 minute. Measure the final temperature of the water using a thermometer. Then measure the final mass of the fuel using a mass scale.
AP Chemistry P2 Experiment 2: Formula of a Hydrate 9/24/2013 Purpose: Calculate the percent composition of water in a hydrate and determine the empirical formula of the hydrate. Procedure: 1) Set up ring stand with ring clamp, clay triangle, crucible with lid, and burner. Adjust the height of the ring stand. 2) Dehydrating Procedures: 3. Measure approximately 1 g of Copper(II) Sulfate Hydrate into the crucible and crucible and lid.
Remove balloons with tongs to show the decrease in volume of the gas inside. Demonstrate that the gas remains inside the balloon by warming the balloon back to room temperature. The flashlight can be used to show that some of the components of air have been condensed. An extension of this demonstration is to distill the air following the procedure by Switzer (J. Chem. Educ.
Introduction In this lab I observed the burning of a candle very closely. I found out that the candle needs oxygen to burn, that it produces carbon dioxide similar to the way that my body produces carbon dioxide, and that a candle produces water as a second waste product. I learned that if I hold an object in the flame it becomes covered with soot which is unburned carbon fuel. Finally, I learned that neither the solid wax, nor the melted wax, nor the wick burns when a candle is lit. In fact, the wax itself is burning as a vapor or gas.
Wait for the bubbling to subside between additions so that the reaction does not overflow the flask. 6. When all the acetic acid has been added, swirl flask or stir for two minutes with a glass stirring rod. 7. When the solution is completely calm, move the flask to a hot plate and heat it to boiling.
Then weigh the crucible without the hydrate after heating. Record both masses. Next, add CoCl2 ∙6H2O and weigh the crucible. Now place the hydrate and crucible on the hot plate. Observe the color change while it is being heated.
Using a Buchner funnel, a hose, and a suction flask we created a vacuum filtrator which we used to help remove the remaining liquid on the copper so that we may make a more precise measurement of the mass of the remaining copper. Our final mass of copper was .7951 grams. Results and Discussion: Initial Mass of Cu: .25 grams When we mixed the 5 ml of 6 molar HNO3 the copper had disappeared, indicating it had been used in the reaction. The copper had undergone a single replacement reaction and a decomposition reaction. Initial equation: Cu(s) + HNO3(aq) -> Cu(NO3)2(aq) + NO2(g) +H2O(l) Balanced: Cu(s) + 4HNO3(aq) -> Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l) The copper had replaced the Hydrogen in the HNO3 and the NO3 had also broken down into NO2 and O2- allowing the H+ to bond with it and create