Enthalpy Lab Report
Abstract
This lab is performed in order to determine the total energy in a reaction between zinc and hydrochloric acid. The reaction is done twice, once to measure the heat of the reaction and again to determine the work done in the system. This is because Enthalpy equals heat plus work (∆H= ∆E+W). Heat and work can be broken down further into separate components so the equation used in lab is ∆H=mc∆T + PV. Many calculations are used in the lab to find out what cannot be measured directly (ex: volume). After all the calculations were complete it was shown to have a very small percent error.
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 enthalpy (∆H) of said reaction. The purpose of this lab is to calculate the enthalpy of a reaction between zinc and hydrochloric acid through the equation ∆H= ∆E+W, which can also be written as ∆H=mc∆T + PV. m is the mass of what is being measured in grams. C is the heat capacity which is “the amount of heat energy required to raise the temperature of a body a specified amount” in joules per gram degree Celsius (chemistry.about.com). ∆T is the final temperature minus the initial temperature in degrees Celsius. The amount of work done by the reaction is equal to the pressure of the surroundings multiplied by the volume of the system. P is the pressure of the surrounding air. V is the volume taken up by the product (H₂).
Experimental Method: The reaction of zinc and hydrochloric acid can be found through a series of ratios that converts grams of zinc into what will fully react with hydrochloric acid is needed. This is found by using the balanced chemical equation Zn+2HCl->H2+ZnC₂. By using this reaction equation in conjunction with the mass of zinc, the mass of hydrochloric acid can be determined with mole ratios. The temperature of the room is then recorded as the initial temperature for both the zinc and the hydrochloric
This lab is performed in order to determine the total energy in a reaction between zinc and hydrochloric acid. The reaction is done twice, once to measure the heat of the reaction and again to determine the work done in the system. This is because Enthalpy equals heat plus work (∆H= ∆E+W). Heat and work can be broken down further into separate components so the equation used in lab is ∆H=mc∆T + PV. Many calculations are used in the lab to find out what cannot be measured directly (ex: volume). After all the calculations were complete it was shown to have a very small percent error.
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 enthalpy (∆H) of said reaction. The purpose of this lab is to calculate the enthalpy of a reaction between zinc and hydrochloric acid through the equation ∆H= ∆E+W, which can also be written as ∆H=mc∆T + PV. m is the mass of what is being measured in grams. C is the heat capacity which is “the amount of heat energy required to raise the temperature of a body a specified amount” in joules per gram degree Celsius (chemistry.about.com). ∆T is the final temperature minus the initial temperature in degrees Celsius. The amount of work done by the reaction is equal to the pressure of the surroundings multiplied by the volume of the system. P is the pressure of the surrounding air. V is the volume taken up by the product (H₂).
Experimental Method: The reaction of zinc and hydrochloric acid can be found through a series of ratios that converts grams of zinc into what will fully react with hydrochloric acid is needed. This is found by using the balanced chemical equation Zn+2HCl->H2+ZnC₂. By using this reaction equation in conjunction with the mass of zinc, the mass of hydrochloric acid can be determined with mole ratios. The temperature of the room is then recorded as the initial temperature for both the zinc and the hydrochloric
References: "Chemistry Definitions." Chemistry - Periodic Table, Chemistry Projects, and Chemistry Homework Help. Web. 28 Nov. 2011. . "Enthalpy." Department of Chemistry at Texas A&M University. Web. 29 Nov. 2011. .