100 Degree Celsius Water To Room Temperature

The dependent variable in the experiment was the temperature and energy absorbed by the water.

3. The third source of error is not knowing for how long to exactly heat the magnesium until it no longer ignites and forms into a white powder of magnesium oxide. In the procedure, it is stated for how long to heat the magnesium, but as the lab went on, it is realized that the magnesium had to be heated for a longer time. This could mean that the magnesium could have been not heated long enough or too little for it to be fully converted into the magnesium oxide product effecting the final results of the lab by having smaller

In order to measure the heats of reactions, add the reactants into the calorimeter and measure the difference between the initial and final temperature. The temperature difference helps us calculate the heat released or absorbed by the reaction. The equation for calorimetry is q=mc(ΔT). ΔT is the temperature change, m is the mass, c is the specific heat capacity of the solution, and q is the heat transfer. Given that the experiment is operated under constant pressure in the lab, the temperature change is due to the enthalpy of the reaction, therefore the heat of the reaction can be calculated.

    In this lab, a calorimeter was used to find the enthalpy of reaction for two reactions, the first was between magnesium and 1 molar hydrochloric acid, and the second was between magnesium oxide and 1 molar hydrochloric acid. After the enthalpy for both of these were found, Hess’ law was used to find the molar enthalpy of combustion of magnesium, using the enthalpies for the two previous reactions and the enthalpy of formation for water. The enthalpy of reaction for the magnesium + hydrochloric acid reaction was found to be -812.76 kJ. The enthalpy of reaction for the magnesium oxide + hydrochloric acid reaction was found to be -111.06 kJ. These two enthalpies and the enthalpy of formation for water were manipulated and added together using Hess’s law to get the molar enthalpy of combustion of magnesium. It was found that the molar enthalpy of combustion of magnesium was -987.5 kJ/mol. The accepted enthalpy was -601.6 kJ/mol, which means that there is a percent difference of 64%. This percent difference is very high which indicates that this type of experiment is very inefficient for finding the molar enthalpy of combustion of magnesium. Most likely, a there are many errors in this simple calorimeter experiment that make it inefficient for finding the molar enthalpy of combustion of magnesium.

In this lab experiment our main focus was to get skillful in using tools such as the metric ruler, balances, thermometer, and graduated cylinder to capture measurements of length, mass, temperature and volume. Additionally, this lab helped us to become more familiar with the uncertainty of measurements, as well as becoming efficient with rounding our measurements to the correct numbers of significant figures. Our results are measured consistently with rounding to the closest answer we could possibly acquire as the data can tell you.

The first part of this lab’s objective was to find the calorimeter constant using DI water. We accomplished this by first checking the temperature and then adding 20 mL of cold DI water into our calorimeter. Next we collected 20 more mL of DI water in a 50 mL beaker and placed it onto a burner in order to heat it. We removed the beaker once the temperature of the water reached 60º celsius. After we removed the water, we poured it into the calorimeter with the cold water and took the temperature. The temperature ended up being 37º celsius. From this information we were able to calculate for qhot, qcold, qcal and Ccal. To be as accurate as possible we conducted this same test three more times and used the averages from all Ccal calculations as the final Ccal.

Because it is dangerous to burn magnesium, it is not possible to directly record heat change. Our lab team suggests an indirect way of determining the heat of combustion for magnesium. To accomplish this, we need to perform two separate trials. One uses a solid (powder) version of MgO, while the other uses Mg ribbon. With the results from these, we can use Hess’ Law to determine q=∆H. This provides both a safe and successful way of indirectly determining the heat of combustion for magnesium.

The purpose of this lab is to determine the percent of water in a hydrate. I learned that copper sulfate hydrate is blue on its own but, when heat is added it will change color to a dim gray. One error that may have occurred is the failure to zero the scale which would ultimately change the math in the equation. Another error that may occur in this lab is forgetting to wait the recommended amount of cooling time, which would change the mass of the elements. At the end if the lab when we look to the questions, it is also beneficial to look at the notes for the lab because it helps with the setup of equations that may come with the questions, also it is very helpful

Therefore, qmetal = -qwater. The mass of the sample was recorded from the baggie. The temperature change

We will be using 6 different fuels to heat up 100ml of water, and find out the changes of the temperature. We will measure the temperatures of the water before and after the experiment. We will burn heat the water for exactly 2 minutes, and check the changes in temperature. The change in temperature will allow us to work out the energy given off the fuel by using this formula:

3. Calculate the total heat released in each reaction, assuming that the specific heat of the solution is the same as for pure water (4.18J/gK). Use q=mcΔT. Show work here and record your answer in Data Table 2.

In this experiment, we performed chemical reactions in a calorimeter to determine the amount of heat produced by heat (temperature) change. In part one of the experiment, we set up our calorimeter as shown in a diagram given. We connected it to our computer for data collection with the temperature probe. To start data collection, we first mixed 15 mL of 1.0 M HCl solution and 20 mL distilled water and recorded the initial temperature. Next, we added 25 mL of 1.0 M NaOH solution to the mixture. To homogenize the mixture, mix gently until the maximum temperature has been reached and record that as the final temperature. Next, we needed to calculate our results for the heat of reaction ΔH in units of Joules. In a chart, we recorded the total mass of our mixture and determined the number of moles

Part A of this experiment was essential because it set up the chemical solutions needed for Part B and C, as well as the equipment that was used throughout the entire experiment. To start, a 300mL solution of 2M HCl was prepared with the stock solution of 6M HCl. For this, 200mL of deionized water was collected and added it to a 400mL beaker. Then, 100mL of 6M HCl was added to the water. Next, 150mL of 2M NaOH was prepped by 50mL deionized water combined with 100mL 3M NaOH in another 400mL beaker. Before the experiment began, the LabQuest was plugged in and turned on. The temperature probe was then plugged into channel one of the LabQuest. In order to fully set up the LabQuest, the interval was changed to 15 s/sample.

Place the calorimeter in the shield. Measure the temperature of water in the calorimeter. Check the thermometer reading after 2 - 3 minutes.

Rinse the thermometer with tap water and dry. Insert the thermometer in the calorimeters top and measure the initial temperature of the base 28.2∘C