Enthalpy Change of Combustion Experiment

2) (0.3 g NaBH4 x 154.16 g ethyl vanillyl alcohol) / (37.83 g NaBH4) = 1.223 g ethyl vanillyl alcohol

The purpose of this experiment was to determine the heat of combustion of magnesium using Hess’s Law. This experiment was significant in understanding how heat can be absorbed or released and how using a calorimeter can help measure heat flow. Part A of this experiment used two dry Styrofoam cups and a magnetic stir bar to create a calorimeter. Temperature was measured and distilled water was placed into the calorimeter, and after 2-3 minutes, boiling water was mixed in. The computer continued to take data for 6 minutes and then the calorimeter was reweighed and emptied. Part B and C were the same, only using two different types of magnesium (magnesium ribbon (part B) and magnesium oxide (part c)) instead of the boiling water. The major result is that all of these experiments release energy because the calorimeter temperature increased.

An Investigation into the Enthalpies of the Combustion of Alcohols = == == == ==

The aim of the lab was to determine the molar enthalpy of the combustion of magnesium using values from reactions between magnesium and magnesium oxide in hydrochloric acid solution using Hess’s Law and Calorimetry. The calculated molar enthalpy of the combustion of magnesium from this experiment was -633.6KJ/mol. This result was found by putting a known mass of magnesium metal as well as magnesium oxide powder into a concentration of HCl. Firstly, a styrofoam calorimeter was used to determine the standard reaction temperature of magnesium metal and hydrochloric acid. Next, calorimetry was again used to determine the reaction temperature of magnesium oxide and hydrochloric acid. Lastly, the thermochemical equations for the two

    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.

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.

In part I, the ΔH of each individual reaction was obtained by performing each reaction inside a calorimeter. Temperature probes were inserted in the calorimeter and ΔT was measured. By using the equation q = Msol’n x Cp x ΔT + Ccal x ΔT, the heat absorbed by the surroundings, q, was obtained for each reaction. The negatives of these values, or heat released by the

where Cc is the concentration of free carbon. The remained molecules of natural gas are mixed to reducing gas and move to the top of furnace. Because the amount of methane in comparison with the amount of reducing gas is very small (about 1%), the changes in temperature profile for both of gas and solid flow are very small. It will be calculated as a function of heat of reaction.

It was desired to compare a theoretical value of enthalpy of combustion to a literature value. To do this, the theoretical value was calculated using a literature value for the heat of sublimation of naphthalene, the heat of vaporization of water and average bond energies, given in Table 1 of the lab packet.1 Equations (1) and (5) were used to calculate the theoretical enthalpy of combustion of gaseous naphthalene, where n was the number of moles, m was the number of bonds, and ΔH was the average bond energy:

Table 1 summarises the results from the primary data collected. Figure 5 indicates there is a linear relationship between the molecular weight and heats of combustion. Figure 6 compares the heats of combustion of the primary data, the accepted values and values calculated from the bond dissociation energy. Figure 6 compares the heat of combustion values identified by the bond dissociation energy calculations, the accepted values and experimental values when 80g of water is heated by 10°C. The bond dissociation energies do not take into account the hydrogen bonding and the accurate energy required to change the tested alkanol from an aqueous state to a gaseous state. The experiments done to identify the accepted values were conducted

When the fuels are burnt, energy is given off. I will be calculating the energy given off using the formula above. The specific heat capacity is the energy

Introduction: The theory behind this experiment is the heat of a reaction (∆E) plus the work (W) done by a reaction is equal to

The aim of this experiment was to test the heat of combustion over a period of time, and the energy required to combust alcohols with different carbon chain levels. It was hypothesised that the higher the carbon chain of the alcohol present, the faster the heat of combustion will occur. Meaning more energy will be released for a higher carbon chain. After calculating the results from the experiment it was found that the hypothesis was partially supported. The reasoning for this is as the alcohol that posses a higher carbon chain, generally increased there reaction rates. However there were a few exceptions to this rule.

Purpose: To measure the heats of reaction for three related exothermic reactions and to verify Hess’s Law of Heat Summation.

And copper (II) ions were reduced to copper because it gained electrons and its oxidation number changed from +2 in copper (II) ions to 0 in copper.