Calorimetry Intro
CALORIMETRY
KATHLEEN IVY B. MENDOZA
DEPARTMENT OF CHEMICAL ENGINEERING, COLLEGE OF ENGINEERING
UNIVERSITY OF THE PHILIPPINES DILIMAN, QUEZON CITY
DATE PERFORMED: JANUARY 30, 2015
INSTRUCTOR’S NAME: JACOB NOEL M. INGUITO
INTRODUCTION
Calorimetry, derived from the Latin calor meaning heat, and the Greek metry meaning to measure, is the science of measuring the amount of heat, q.1 The amount of heat is absorbed (endothermic where qrxn >0) or released (exothermic where qrxn <0) by the system with respect to its surroundings is the change in enthalpy, ΔH. This is quantified in calorimetry by a device called calorimeter.
There are two types of calorimeters. The first type is the called the bomb calorimeter. It is ideally suited for measuring the heat evolved in a combustion reaction. Because the bomb confines the reaction mixture to a fixed volume, the reaction is said to occur at constant volume.
The other, much simpler and the calorimeter you are most likely to encounter is the coffee-cup calorimeter. This uses Styrofoam, a good heat insulator, for very minimal transfer of heat between the cup and the surrounding air. The reaction in the calorimeter occurs under constant pressure of the atmosphere.2
Both are isolated—it does not permit any exchange of matter and energy—and adiabatic—there is no heat transfer between the system and its surroundings (qsystem=0). Thus, this equation occurs. qsystem = qcal + qrxn = 0 qcal = -qrxn In this experiment, a styrofoam calorimeter was utilized to obtain the changes in enthalpy of the various reactions. Before being able to do this, however, the heat capacity of the calorimeter, Ccal, must be determined. This is accomplished by the neutralization of a strong acid and strong base wherein the change of enthalpy of the reaction is already known.3 Having found the Ccal, the calculations for the enthalpy changes and the specific heat of a metal can be performed.
qcal = CcalΔT (2) qrxn = nΔHrxn (3)
Substituting Equation 2
KATHLEEN IVY B. MENDOZA
DEPARTMENT OF CHEMICAL ENGINEERING, COLLEGE OF ENGINEERING
UNIVERSITY OF THE PHILIPPINES DILIMAN, QUEZON CITY
DATE PERFORMED: JANUARY 30, 2015
INSTRUCTOR’S NAME: JACOB NOEL M. INGUITO
INTRODUCTION
Calorimetry, derived from the Latin calor meaning heat, and the Greek metry meaning to measure, is the science of measuring the amount of heat, q.1 The amount of heat is absorbed (endothermic where qrxn >0) or released (exothermic where qrxn <0) by the system with respect to its surroundings is the change in enthalpy, ΔH. This is quantified in calorimetry by a device called calorimeter.
There are two types of calorimeters. The first type is the called the bomb calorimeter. It is ideally suited for measuring the heat evolved in a combustion reaction. Because the bomb confines the reaction mixture to a fixed volume, the reaction is said to occur at constant volume.
The other, much simpler and the calorimeter you are most likely to encounter is the coffee-cup calorimeter. This uses Styrofoam, a good heat insulator, for very minimal transfer of heat between the cup and the surrounding air. The reaction in the calorimeter occurs under constant pressure of the atmosphere.2
Both are isolated—it does not permit any exchange of matter and energy—and adiabatic—there is no heat transfer between the system and its surroundings (qsystem=0). Thus, this equation occurs. qsystem = qcal + qrxn = 0 qcal = -qrxn In this experiment, a styrofoam calorimeter was utilized to obtain the changes in enthalpy of the various reactions. Before being able to do this, however, the heat capacity of the calorimeter, Ccal, must be determined. This is accomplished by the neutralization of a strong acid and strong base wherein the change of enthalpy of the reaction is already known.3 Having found the Ccal, the calculations for the enthalpy changes and the specific heat of a metal can be performed.
qcal = CcalΔT (2) qrxn = nΔHrxn (3)
Substituting Equation 2
References: (accessed February 3, 2015) [2] Petrucci, R.H., Harwood, W.S., Herring, F.G [3] Beran, J.A. Laboratory Manual for Principles of General Chemistry, 9th ed.; John Wiley & Sons, New Jersey, 2011; pp.287-289 Brown, T.E., LeMay, E.H., Bursten, B.E (accessed February 3, 2015) [4] Davidson College Chemistry Resources http://www.chm.davidson.edu/ (accessed February 3, 2015)