Lab 1 report chem 12400 PDF

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PHSC 12400 Lab 1 Report

Chemistry over fire or ice? This experiment focused on endothermic and exothermic reactions. Endothermic reactions are those that require energy, like breaking chemical bonds, and exothermic reactions, like forming chemical bonds, release energy. Since energy is conserved, change in energy can be measured through changes in temperature levels. Exothermic reactions lead to an increase of temperature by releasing heat and vice versa for endothermic reactions. These reactions are very important to life, like photosynthesis for example. In this series of endothermic reactions, plants absorb energy from sunlight to turn carbon dioxide and water into oxygen and glucose, where the energy is stored until needed. Exothermic reactions then are used to release this energy. The objective of this lab was to understand endothermic and exothermic reactions like this on a smaller scale, using temperature changes to try to classify reactions as endo or exothermic. For this lab, the first step was to measure out water (anywhere from 5 to 10 mL) into a graduated cylinder and transfer that into a test tube after writing down the amount. Then, we used a thermometer to record the initial temperature of the water. Next step was to weigh out approximately 0.5 g of solid ammonium nitrate using a benchtop balance and add that to the test tube. After swirling and dissolving the solid, we recorded the final temperature of the test tube. We then repeated all the steps again with calcium chloride instead of the solid ammonium nitrate.

Ammonium nitrate

Calcium chloride

Water volume 6.1 mL

Mass

Initial T 0.64g 26.2 deg C

Final T 23.9 deg C

5.3 mL

0.50g 26.4 deg C

27.8 deg C

Delta T 2.3 deg C ± 1.15 1.4 deg C ± 0.7

Observations Felt cooler after mixing

Endo or Exothermic? Endothermic reaction

Felt warmer after mixing

Exothermic reaction

Chemical reactions: NH4NO3(s) -> NH4+(aq) + NO3- (aq) CaCl2 (s) + 2H2O (l) -> 2HCl (aq) + Ca(OH)2 (s)

The results above clearly prove what we know about exothermic and endothermic reactions and what we expected to find. Inadvertent error was definitely possible, through the measuring of the water used and the weighing of the solids used, as well as the reading of the thermometers. Our data makes sense as is, but if we were doing further calculations or finding Gibbs free energy, errors would be much more significant. Future work could be more trials of this experiment with different temperature levels of water and maybe also with varying amounts of solute. It would also be interesting to try and find the Gibbs free energy value to investigate the spontaneity aspect of these reactions. For the ammonium nitrate reaction, the temperature decreased, as heat was lost during the dissolution of the solid into aqueous ammonium and nitrate. The bonds were broken between ammonium and nitrate. Energy was required to break these chemical bonds in ammonium nitrate once mixed with water, absorbing heat and making the tube feel colder to the touch, which is why it was an endothermic reaction. The calcium

chloride reaction was exothermic because the final temperature increased after the rearranging of the bonds. The bonds between calcium and chloride were broken and then were reformed with the ions in water, releasing energy in the process. Heat was given off in this reaction, so the tube felt slightly warmer after the mixing. Since there was a higher overall change of temperature for the ammonium nitrate reaction, there was probably a greater amount of energy involved than for the calcium chloride reaction, which had a change of only about 1 degree Celsius. An example of an exothermic reaction outside the classroom is water freezing into ice. For water to become its solid form, the particles come close together and more bonds are formed. As discussed above, when bonds are formed, energy is released, making this an exothermic reaction. An endothermic reaction is the reverse reaction of this. Bonds are broken as ice melts into water, so energy is needed for this to occur. Spontaneous endothermic reactions are identified by the value of the Gibbs free energy of a system, which is equal to enthalpy minus the product of the temperature and entropy of the system. If the Gibbs free energy is negative, the reaction will occur spontaneously. Endothermic reactions generally require energy to run, but not all of them. Thus, whether the endothermic reaction occurs spontaneously or not depends on the temperature, enthalpy, and entropy values of the system and how they fit into the Gibbs free energy equation.1 If the experiment was repeated with sodium chloride, a similar result to the ammonium nitrate reaction would occur. The bonds between the sodium and chloride

!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!! 1 !Gibbs!Free!Energy,!Purdue&University,! https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch21/gibbs.php

ions would break and the Na+ and Cl- ions would dissociate into the water. This would require energy for the chemical bonds to break, absorbing heat as an endothermic reaction. As for sucrose, since it is made with covalent bonds, it is much stronger than table salt. The molecule itself won’t come apart, but the whole sucrose molecules will separate from one another when it dissolves. These molecules will mix with the water because both are polar substances. The sugar molecules will form intermolecular bonds with the water molecules, giving off energy as an exothermic reaction. Overall, this lab proved that breaking chemical bonds requires energy by decreasing the final temperature, making it an endothermic reaction. Similarly, it showed that forming chemical bonds releases energy by increasing the final temperature, making it an exothermic reaction. I learned about how to predict whether a reaction can be endothermic or exothermic based on chemical equations and my knowledge of polarity and bonding. This is applicable to every day concepts like cold packs, which are made with solid ammonium and water. You just crush the contents and mix the contents, which kick starts the endothermic reaction. The pack absorbs heat from its surroundings, in this case “freezing” the painful area of one’s body.2

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