Calorimetry

11 Key excerpts on "Calorimetry"

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  Chemistry

  • Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
  • 2015(Publication Date)
  • Openstax

    (Publisher)

  Knowledge of the heat capacity of the surroundings, and careful measurements of the masses of the system and surroundings and their temperatures before and after the process allows one to calculate the heat transferred as described in this section. A calorimeter is a device used to measure the amount of heat involved in a chemical or physical process. For example, when an exothermic reaction occurs in solution in a calorimeter, the heat produced by the reaction is absorbed by the solution, which increases its temperature. When an endothermic reaction occurs, the heat required is absorbed from the thermal energy of the solution, which decreases its temperature (Figure 5.11). The temperature change, along with the specific heat and mass of the solution, can then be used to calculate the amount of heat involved in either case. 238 Chapter 5 | Thermochemistry This OpenStax book is available for free at http://cnx.org/content/col11760/1.9 Figure 5.11 In a calorimetric determination, either (a) an exothermic process occurs and heat, q, is negative, indicating that thermal energy is transferred from the system to its surroundings, or (b) an endothermic process occurs and heat, q, is positive, indicating that thermal energy is transferred from the surroundings to the system. Scientists use well-insulated calorimeters that all but prevent the transfer of heat between the calorimeter and its environment. This enables the accurate determination of the heat involved in chemical processes, the energy content of foods, and so on. General chemistry students often use simple calorimeters constructed from polystyrene cups (Figure 5.12). These easy-to-use “coffee cup” calorimeters allow more heat exchange with their surroundings, and therefore produce less accurate energy values. Chapter 5 | Thermochemistry 239 Figure 5.12 A simple calorimeter can be constructed from two polystyrene cups. A thermometer and stirrer extend through the cover into the reaction mixture.

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  Thermal Analysis

  • Bernhard Wunderlich(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  C H A P T E R 5 Calorimetry Calorimetry represents the effort to measure heat (caloric, see Fig. 1.1) in any of its manifestations. 1 This attempt to measure heat directly distinguishes the present discussion from Chapters 3 and 4, in which the measurement of temperature did not lead to quantitative, caloric information. There is, however, no heat meter, meaning there is no instrument which allows one to find the heat content of a system directly, as has been mentioned already in Sect. 4.1. The measurement of heat must always be made in steps and summed from a chosen initial state. The most common reference state is that of the chemical elements, stable at 298.15 Κ [Δ/ff (298) = 0; see Fig. 2.14]. 5.1 Principles and History The three common ways of measuring heat are listed at the top of Fig. 5.1. First, the change of temperature in a known system can be observed and related to the flow of heat into the system. It is also possible, using the second method, to follow a change of state, such as the melting of a known system, and determine the accompanying flow of heat from the amount of material transformed in the known system. Finally, in method three, the conversion to heat of known amounts of chemical, electrical, or mechanical energy can be used to duplicate (or compensate) the flow of heat, and thus measure heat by comparison. The prime difficulty of all calorimetric measurements is the fact that heat cannot be contained. There is no known perfect insulator for heat. During the time one performs the measurement, there are continuous losses. All Calorimetry is thus beset by efforts to make corrections for heat loss. Matter always contains thermal energy, and this thermal energy is constantly exchanged. Even if there were a perfect vacuum surrounding the system under investigation, heat would be lost and gained by radiation. 219 220 Thermal Analyste PRINCIPLES AND HISTORY In Calorimetry, heat measurements involve: 1.

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  Science of Heat and Thermophysical Studies

  A Generalized Approach to Thermal Analysis

  • Jaroslav Sestak(Author)
  • 2005(Publication Date)
  • Elsevier Science

    (Publisher)

  Chapter 12

  THERMOMETRY AND Calorimetry

  12.1. Heat determination by Calorimetry

  Various thermometric assessments have been in the center of retailored techniques used to detect a wide variety of heat effects and properties. The traditional operation aims to measure, for example, heat capacities, total enthalpy changes, transitions and phase change heats, heats of adsorption, solution, mixing, and chemical reactions. The measured data can be used in a variety of clever ways to determine other quantities. Special role was executed by methods associated with enough adequate temperature measurements, which reveals an extensive history coming back to the first use of the word ‘calorimeter’ introduced by the work Wilcke and later used by Laplace, Lavoisier as already discussed in the previous Chapter 4 .

  Calorimetry is a direct and often the only way of obtaining the data in the area of thermal properties of materials, today especially aimed to higher temperatures. Detailed descriptions are available in various books [3 , 9 , 590 –596 ] and reviews [597 –599 ]. Although the measurements of heat changes is common to all calorimeters, they defer in how heat exchanges are actually detected, how the temperature changes during the process of making a measurement are determined, how the changes that cause heat effects to occur are initiated, what materials of construction are used, what temperature and pressure ranges of operation are employed, and so on. We are not going to describe herewith the individual peculiarities of instrumentation as we merely focus our attention to a brief methodical classification.

  If the calorimeter is viewed as a certain ‘black box’ [3] , whose input information are thermal processes and the output figures are the changes in temperature or functions derived from these changes. The result of the measurement is an integral change whose temperature dependence is complicated by the specific character of the given calorimeter and of the measuring conditions. The dependence between the studied and measured quantity is given by a set of differential equations, which are difficult to solve in the general case. For this reason, most calorimetric measurements are based on calibration. A purely mathematical solution is the domain of a few theoretical schools [3 , 594 , 596

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  Chemistry

  Principles and Reactions

  • William Masterton, Cecile Hurley(Authors)
  • 2020(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  187 8 ▼ Thermochemistry Chapter Outline 8-1 Principles of Heat Flow 8-2 Measurement of Heat Flow; Calorimetry 8-3 Enthalpy 8-4 Thermochemical Equations 8-5 Enthalpies of Formation 8-6 Bond Enthalpy 8-7 The First Law of Thermodynamics ▼ T his chapter deals with energy and heat, two terms used widely by both the gen-eral public and scientists. Energy, in the vernacular, is equated with pep and vitality. Heat conjures images of blast furnaces and sweltering summer days. Scientifically, these terms have quite different meanings. Energy can be defined as the capacity to do work. Heat is a particular form of energy that is transferred from a body at a high temperature to one at a lower temperature when they are brought into contact with each other. Two centuries ago, heat was believed to be a material fluid (caloric); we still use the phrase “heat flow” to refer to heat transfer or to heat effects in general. Thermochemistry refers to the study of the heat flow that accompanies chemical reactions. Our discussion of this subject will focus on ■ the basic principles of heat flow (Section 8-1). ■ the experimental measurement of the magnitude and direction of heat flow, known as Calorimetry (Section 8-2). ■ the concept of enthalpy, H (heat content) and enthalpy change , D H (Section 8-3). ■ the calculation of D H for reactions, using thermochemical equations (Section 8-4) and enthalpies of formation (Section 8-5). ■ heat effects in the breaking and formation of covalent bonds (Section 8-6). ■ the relation between heat and other forms of energy, as expressed by the first law of thermodynamics (Section 8-7). Scala/Art Resource, NY The candle flame gives off heat, melting the candle wax. Wax melting is a phase change from solid to liquid and an endothermic reaction. Some say the world will end in fire, Some say in ice. From what I’ve tasted of desire I hold with those who favor fire. —ROBERT FROST Fire and Ice Copyright 2016 Cengage Learning.

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  Pharmaceutical Isothermal Calorimetry

  • Simon Gaisford, Michael A. A. O'Neill(Authors)
  • 2006(Publication Date)
  • CRC Press

    (Publisher)

  here. It is, however, important to know the basic designs and operating principles underpinning all modern calorimeters in order to understand the origin of calori-metric signals and to draw comparison between them. Measuring Principles There are only three methods by which heat can be experimentally measured: 1. Measurement of the power required to maintain isothermal conditions in a calorimeter, the power being supplied by an electronic temperature controller in direct contact with the calorimeter (power compensation Calorimetry). 2. Measurement of a temperature change in a system which is then multiplied by an experimentally determined cell constant (adiabatic Calorimetry). 3. Measurement of a temperature difference across a path of fixed thermal conductivity which is then multiplied by an experimentally determined cell constant (heat conduction Calorimetry). Note that all calorimetric measurements therefore require a minimum of two experiments (one for measurement and one for calibration), although further measurements may be needed for blank corrections (such as the determi-nation of a baseline or the correction for dilution enthalpies in a titration experi-ment). A discussion of the need for, and methods of, calibration (including chemical test reactions) is the basis of Chapter 2. Power Compensation Calorimeters In power compensation Calorimetry, an electrical element is used either to add heat or remove heat from the calorimetric vessel as the sample reacts, maintain-ing the sample and vessel at a given temperature. The power output from the sample is thus the inverse of the power supplied by the element. In order to be able to heat and cool, the element is usually based on the Peltier principle. A typical application of this type of Calorimetry is power compensation DSC. Adiabatic Calorimetry In an ideal adiabatic calorimeter, there is no heat exchange between the calori-metric vessel and its surroundings.

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  Physics for O.N.C. Courses

  • R.A. Edwards(Author)
  • 2014(Publication Date)
  • Pergamon

    (Publisher)

  7.2 The Experimental Determination of Specific Heats. Calorimetry

  Any vessel designed to contain materials for the purpose of measuring heat quantities developed or exchanged by these materials is called a calorimeter. In its simplest form a calorimeter consists of a cylindrical can made usually of copper, which is a good conductor of heat and so may be assumed to undergo the same change in temperature as the contents. A so-called “continuous flow” calorimeter was used by Callendar and Barnes to determine the specific heat of water with great accuracy over the range of temperature from 0° to 100°C. The essential features of their apparatus is illustrated in Fig. 7.1 . Water from a constant head flows at a uniform rate over an electrically heated element situated along the length of the inner glass tube of the calorimeter. The vacuum jacket and water jacket which surround the inner tube are also of glass and serve to minimise heat loss to the surroundings.

  FIG. 7.1

  A steady current is maintained in the heating element and after some time a steady state is set up in which the readings of the two thermometers no longer change. The temperature θ2 of the water as it leaves the calorimeter is higher than its initial temperature θ1 , these two temperatures being recorded by thermometers T 2 and T 1 respectively. Since, in the steady state, these temperatures remain constant with time, they may be measured very accurately (e.g. using platinum-resistance thermometers, as did Callendar and Barnes) and so θ2 may be quite close to θ1 without loss of accuracy. In the steady state, the energy supplied electrically is equal to the heat carried off by the flowing water in the same time, so that

  (7.5)

  where s is the specific heat of the water in J g−1 degC−1 or in J kg−1 degC−1 if m is the mass of water in grams or kilograms, respectively, flowing through in t sec, V is the p.d. in volts across the heater and I the current in amps through it (held constant throughout). The (small) loss of heat to the surroundings is represented by h

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  Principles of Thermodynamics

  • Jean-Philippe Ansermet, Sylvain D. Brechet(Authors)
  • 2019(Publication Date)
  • Cambridge University Press

    (Publisher)

  Part II PHENOMENOLOGY 5 Calorimetry Benoît Paul Emile Clapeyron, 1799–1864 B. P. E. Clapeyron was a Frenchman who today is known for his representation of thermal cycles by (p, V) diagrams. He found a relation between the slope of a p (T) line in a phase diagram and the corresponding latent heat. 5.1 Historical Introduction In this chapter, we will introduce the notion of specific heat, also known as heat capacity. Antoine Laurent de Lavoisier and Pierre-Simon de Laplace designed an experiment that allowed them to quantify heat phenomena [42]. They built an ice calorimeter (Fig. 5.1) 103 104 Calorimetry Figure 5.1 Calorimeter of Lavoisier and Laplace. When a thermodynamic process takes place inside the inner chamber, heat is released which makes the ice melt. The heat transfer is quantified by the amount of water collected. which they used to determine the heat released in a chemical reaction. The reaction took place inside a sphere of ice kept at 0 ◦ C. From the amount of ice that melted in that process, they measured the heat of the reaction, also called reaction enthalpy. The heat transfer that occurred when a set amount of gas passed through the ice was measured by the amount of ice that melted. They also recorded the temperature of the gas at the inlet and the outlet. The heat value thus obtained, divided by the temperature drop, yielded the heat capacity or the specific heat of the gas. These famous scientists also used their calorimeter to analyse the breathing of animals and concluded from their observations that breathing must be a combustion process. Experimentally, it is observed that the specific heat of any substance decreases to zero as the temperature tends to zero. As we will see in this chapter, this general property can be derived from what is known as the third law of thermodynamics.

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  Thermodynamics of Natural Systems

  Theory and Applications in Geochemistry and Environmental Science

  • Greg Anderson(Author)
  • 2017(Publication Date)
  • Cambridge University Press

    (Publisher)

  This gives rise to another kind of Calorimetry, cryogenic , or low-temperature Calorimetry. 5.4.3 Cryogenic Calorimetry A cryogenic calorimeter ( Figure 5.5 ) is an apparatus designed for the determination of heat capacities at very low temperatures. The procedure is to cool the sample down to a temperature within a few degrees of absolute zero (a temperature of absolute zero itself Figure 5.5 A cryogenic or low-temperature calorimeter. The sample container can be raised by the rotary winch so as to be in contact with the liquid helium reservoir for cooling to 4.2 K, or lowered into the vacuum for heating. The re-entrant well in the sample container contains a heating coil. (Simplified from Robie and Hemingway (1972).) 121 5.5 The Problem Resolved (a) C P , J mol –1 K –1 C P / T (b) Figure 5.6 (a) Measured heat capacity of muscovite as a function of temperature (Robie et al. , 1976). (b) C P / T vs. T for the same data. Integration gives the shaded area under the curve, which is equal to the entropy at the upper limit of integration, in this case, S ◦ 298.15 = 287.7 J mol − 1 . is actually impossible to achieve, a fact actually implicit in the third law), introduce a known quantity of heat using an electrical heating coil, and observe the resulting increase in temperature (usually a few degrees). The quantity of heat is equal to H , and this divided by the temperature difference gives an approximate value of C P at the midpoint of the temperature range. Corrections are then made to compensate for heat leaks, for the heat absorbed by the calorimeter, and to get exact C P values from the approximate ones. The integration of C P / T values to obtain the entropy at 298.15 K is illustrated in Figure 5.6 . A much more detailed description of the calorimeter and its operation is in Robie (1987). 5.5 The Problem Resolved ...............................................................................

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  Research Methodologies in Human Diabetes. Part 2

  • C. E. Mogensen, E. Standl, C. E. Mogensen, E. Standl(Authors)
  • 2020(Publication Date)
  • De Gruyter

    (Publisher)

  Direct Calorimetry measures the total heat lost from an organism to the environment, whereas indirect Calorimetry measures respiratory gas exchange, from which an inference is made on the rate and type of oxidative reactions that release the free energy of metabolic fuels. Thus, the two calorimetric methods provide estimates for the two sides of the energy balance equation that must coincide when measurements are ob-tained at equilibrium and are free of systematic errors. Indeed, when the two techniques are combined, the respective estimates of energy turnover become super-imposable within a half hour [1]. 99 E. Ferrannini Direct Calorimetry The equation upon which direct Calorimetry is based is: M = (R + C + K + E) +S where R, C, K, and E are the radiant, convective, conductive, and evapora-tive heat transfer rates, and S is the rate of body heat storage. For a subject at rest, whose body heat content is constant (S = O), the sum total of all 4 routes of heat dissipation will measure heat production (M). In practice, an isothermal calorimeter is a thermally insulated chamber surrounded by a shell space in which air is maintained at the same temper-ature as that of the chamber itself. In this gradient-free system, air temper-ature and humidity are measured at the chamber inlet and outlet. The body dry and evaporative heat output are then calculated from the temperature and humidity gradient between inlet and outlet and the air flow (corrected for the heat added or substracted by the heat exchanger). In gradient-layer Calorimetry, the body is tightly wrapped in a substance (an epoxy resin) of known thermal conductivity and thickness, and the tem-perature gradient between the inner and outer layers provide a measure of (R + C + K), while E is measured by condensing out the exact amount of water lost by the subject.

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  Cell Chemistry and Physiology: Part III

  • Edward Bittar(Author)
  • 1996(Publication Date)
  • Elsevier Science

    (Publisher)

  Scand. J. Infect Dis. Suppl. 9, 12-16. Marsh, K.N. & O'Hare, P.A.G. (eds.) (1994). In: Experimental Chemical Thermodynamics. Vol. 2. Solution Calorimetry. Blackwell, Oxford. Mills, I., Cvitas, T., Homan, K., Kallay, N., & Kuchitsu, K. (1993). In: Quantities, Units and Symbols in Physical Chemistry. Blackwell, Oxford. Monti, M. (1987). In vitro thermal studies of blood cells. In: Thermal and Energetic Studies of Cellular Biological Systems (James, A.M., ed.), pp. 131-146, Wright, Bristol. Calorimetric Techniques 301 Monti, M., Ikomi-Kumm, J., Ljunggren, L., Lund, U., & Thysell, H. (1993). Medical application of microCalorimetry in human toxicology. A study of blood compatibility of haemodialysis mem-branes. Pure Appl. Chem. 65,1979-1981. Newell, R.D. (1980). The identification and characterization of microorganisms by microCalorimetry. In: Biological MicroCalorimetry (Beezer, A.E., ed.), pp. 163-186, Academic Press, London. Parrish, W.R. & Lewis, E.A. (eds.) (1996). Handbook of Calorimetry. Marcel Dekker, New York. In press. Privalov, P.L. (1980). Heat capacity studies in biology. In: Biological MicroCalorimetry (Beezer, A.E., ed.), pp. 413-451, Academic Press, London. Privalov, P.L. (1989). Thermodynamic problems of protein structure. Ann. Rev. Biophys. Chem. 18, 47-69. Randzio, S. & Suurkuusk, J. (1980). Interpretation of calorimetric thermograms and their dynamic corrections. In: Biological MicroCalorimetry (Beezer, A.E., ed.), pp. 311-341, Academic Press, London. Ross, P.D. & Goldberg, R.N. (1974). A scanning microcalorimeter for thermally induced transitions in solution. Thermochim. Acta 10,143-151. Schon, A. & Wadso, I. (1988). The potential use of microCalorimetry in predictive tests of the action of antineoplastic drugs on mammalian cells. Cytobios 55,33-39. Sturtevant, J.M. (1971). Calorimetry. In: Physical Methods of ChemisU^ (Weissberger, A. & Rossiter, B.W., eds.). Vol. 1, Part V, pp. 347-425, Wiley, New York. Wadso, I.

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  Thermodynamics of Geothermal Fluids

  • Andri Stefánsson, Thomas Driesner, Pascale Bénézeth(Authors)
  • 2018(Publication Date)
  • De Gruyter

    (Publisher)

  The first reported application of flow technology to calorimetric measurements was by Priestley et al. (1965), who used an isoperibol design which measured the tem-perature difference at the outlet of the mixing chamber relative to the mean temperature of the two incom-ing liquids. His instrument was used to measure the enthalpies of formation of a series of metal complexes with EDTA. In the late 1960's, Stoesser and Gill (1967) and Monk and Wadso (1968) reported independent designs for twin-cell heat-flux calorimeters. The twin cell design uses a reference cell, identical to the mixing cell but either without flow, or with flow of the mixed liquid, to cancel non-ideal heat flux behavior. In 1969, Picker et al. published novel designs two heat-of-mixing micro-calo-rimeters. The first was an adiabatic instrument that measured the temperature increment occur-ring during the mixing process relative to the inlet temperature of the two liquids. The second design was for a flow calorimeter capable of running under adiabatic or isothermal conditions to study both liquid- and gas phase reactions. Both instruments were so sensitive that they were able to reach steady states within 1 min, so that the flow rates of the mixing fluids could be varied continuously in order to obtain heats of mixing over the entire composition range. Principles of operation for isothermal calorimeters. The underlying operating principle of isothermal calorimeters is based on heating or cooling the reaction vessel to balance the heat liberated or consumed by the mixing reaction. In order to maintain the reaction zone at a constant temperature the energy output is adjusted with a controlled heater to balance the energy arising from the chemical reaction plus the energy removed by a constant heat-leak path. The enthalpy is directly obtained from the power Af (mW) required to maintain the temperature of the calorimeter constant and the molar flow-rate f n (mol-s -1 ) of the solution.

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