Thermodynamics Lab Report PDF

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Neela Mahabir Thermodynamics Chemistry I Section 05 12/04/2020

Abstract Energy is one of the great subject matters of our time, but what is energy actually? In this simulation I learned the fundamental thermodynamic concepts of enthalpy, entropy and Gibbs free energy. I, also determine the internal energy of a chemical compound by using bomb calorimetry, and was able to even travel inside the calorimeter to see it in action. In this Lab, I was able to define the core thermodynamics concepts of entropy, enthalpy, and free Gibbs energy, and their units. Explain the first and second laws of thermodynamics. Understand and apply the concept of reaction spontaneity. Explain the differences between the enthalpy of combustion, and enthalpy of formation. Understand the relationship between internal energy and enthalpy. Present Hess’s law in connection with performing enthalpy change calculations and present the concepts of exothermic and endothermic reactions

Introduction Thermodynamics is the branch of physics that deals with the relationships between heat and other forms of energy. In particular, it describes how thermal energy is converted to and from other forms of energy and how it affects matter. Gibbs Free Energy, G, can be used to determine if a reaction is spontaneous or not. A negative value of G indicates that a given reaction is spontaneous at the measured conditions and will proceed in the forward direction. G can be calculated using the Gibbs-Helmholtz equation: G = H – TS. In this experiment, H (enthalpy) will be calculated from the temperature change of the reaction. S (entropy) will be calculated using standard entropy values from the textbook.

The reactions used in this lab will be the dissolution (dissociation) of two salts in water. It will be important to differentiate between the system (or reaction) and the surroundings in this experiment. The calorimeter used in this experiment is assumed to be a closed, isolated container that does not lose any heat to the environment. Therefore, all heat exchanges are assumed to take place between the system and the surroundings. The dissolution of each salt is the system; water is the surroundings. It is the temperature change of the water (the surroundings) that will be measured over time.

Abstract:

In this experiment, the heat of formation of magnesium oxide was calculated by measuring the enthalpies of other reactions that could then be added together to obtain the heat of formation of magnesium oxide using Hess’s Law. The two other reactions that were conducted in the experiment, were the reaction between

solid magnesium and hydrochloric acid, as well as the reaction between magnesium oxide and hydrochloric acid. Knowing these two reactions, and the heat of formation of water, the heat of formation of magnesium oxide can be found. The specific heat of an unknown metal was also

found in the lab, through the use of q=mcΔT. Introduction: For this lab, one of the most important concepts that were applied was Hess’s Law. This law states that enthalpy for a reaction is equivalent to the addition of enthalpies from reactions that can be added together to get the

desired reaction. These reactions can be seen as steps of the overall reaction. Another important aspect of the lab was the transfer of heat from the reaction to the surroundings. The equation q=mcΔT can be used in order to determine the amount of heat transferred from the reaction to the surrounding water. This

calculated heat can then be used to determine the specific heat capacity of the unknown metal. Abstract: In this experiment, the heat of formation of magnesium oxide was calculated by measuring the enthalpies of other reactions that could then be added

together to obtain the heat of formation of magnesium oxide using Hess’s Law. The two other reactions that were conducted in the experiment, were the reaction between solid magnesium and hydrochloric acid, as well as the reaction between magnesium oxide and hydrochloric acid.

Knowing these two reactions, and the heat of formation of water, the heat of formation of magnesium oxide can be found. The specific heat of an unknown metal was also found in the lab, through the use of q=mcΔT. Introduction: For this lab, one of the most important concepts

that were applied was Hess’s Law. This law states that enthalpy for a reaction is equivalent to the addition of enthalpies from reactions that can be added together to get the desired reaction. These reactions can be seen as steps of the overall reaction. Another important aspect of the

lab was the transfer of heat from the reaction to the surroundings. The equation q=mcΔT can be used in order to determine the amount of heat transferred from the reaction to the surrounding water. This calculated heat can then be used to determine the specific heat capacity of the unknown metal.

For this lab, one of the most important concepts that were applied was Hess’s Law. This law states that enthalpy for a reaction is equivalent to the addition of enthalpies from reactions that can be added together to get the desired reaction. These reactions can be seen as steps of the

overall reaction. Another important aspect of the lab was the transfer of heat from the reaction to the surroundings. The equation q=mcΔT can be used in order to determine the amount of heat transferred from the reaction to the surrounding water. This calculated heat can then be used to

determine the specific heat capacity of the unknown metal. It had been recognized by the 18th century that the amount of heat, Q, required to change the temperature of a system is proportional to the mass, m, of the system and to the temperature change, ∆T. This proportion is what we now know as “specific heat” and is often abbreviated by c. This value is valid for heat flow into or out of the system. This lab will help us see how to determine the specific heat of a known substance.

Method

1. Place the cup into a beaker. Use a graduated cylinder to measure deionized water and add to the blue plastic cup inside the cups. Record the exact volume of water used.

2. Tare out (zero) the mass of the weighing cup. Place a thermometer in the DI water. This will be your initial temperature. 3. Add the sample in the calorimeter and replace the lid. Record the temperature of the mixture every 10 seconds. 4. Be sure to record your observations of the appearance of the salt solution in the calorimeter at the end of two minutes. .

Materials requires: Calorimeter

Goggles

Lab Coat

Thermometer

Cup/lid

Samples

Analytical Scale

Results and Discussion In the experiment, the recorded temperature change for the sample reaction. Using this information and the known specific heat capacity of water, the heat of reaction can be calculated using q=mcT. The calculated heat of the reaction turns out to be 4.18 kJ. This means that the water absorbed 4.18 kJ of energy from the reaction. The data shows a clear trend that as the temperature decreases, the concentration of the tetraborate ion decreases, which follows that the value of Ksp decreases as temperature decreases (Table 1), which thus follows that the value of ln(Ksp) decreases as temperature decreases. This value means that the changes in 1/T(K) explains 99.15% of the variations in ln(Ksp). This is good because it shows that there are not any influential points that can drastically change the line of best fit, which could also drastically change the values of the slope and the intercept. The results show the trend that as temperature decreases, the value of the equilibrium constant also decreases. This satisfies the first goal of this experiment in studying the temperature-dependence of the equilibrium constant. The second goal of this experiment was to determine the values of the enthalpy and entropy of dissolution. The

signs of the enthalpy and entropy of dissolution are both positive. It makes sense that the ΔHº of solution is positive, meaning that the reaction is endothermic, because the mixture was heated and therefore the molecules needed to absorb the heat in order to turn into a solution. This also helps to understand why the sign of ΔSº is positive. When the signs of both ΔHº and ΔSº are positive, the reaction will be spontaneous at high temperatures (i.e. ΔGº < 0).

Figure 1.1

Figure 1.2

Table 1

Conclusion In conclusion, the first and second laws of thermodynamics and the core concepts of enthalpy, entropy and Gibbs free energy are introduced in this simulation. I was able to play around with the energy levels of reactants and products on a virtual energy surface to learn about endothermic and exothermic reactions. The concept of reaction spontaneity is linked to the concept of Gibbs free energy and its temperature dependence is explored in an interactive game. I had access to a state-of-the-art bomb calorimeter and can travel inside to see it in action, in order to really be able to understand how it works. From here the concept of chemical bond energy is linked to the thermodynamic calculations of enthalpy on the calorimeter output.

Reference:

Thermodynamics. (n.d.). Retrieved December 04, 2020, from https://www.britannica.com/science/thermodynamics

What is Thermodynamics? (n.d.). Retrieved December 04, 2020, from https://www.universalclass.com/articles/science/what-is-thermodynamics.htm...