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Chapter 6: Problem 30

Define calorimetry and describe two commonly used calorimeters. In a calorimetric measurement, why is it important that we know the heat capacity of the calorimeter? How is this value determined?

Short Answer

Expert verified
Calorimetry is the measurement of heat transfer in a physical or chemical process. Two common calorimeters are the coffee cup and the bomb calorimeter. Knowing the calorimeter's heat capacity is important because it absorbs some heat which affects the overall measurement of the heat transfer. This heat capacity is usually determined by using a substance with known specific heat capacity or a reaction with a known heat of reaction.

Step by step solution

01

Definition of Calorimetry

Calorimetry is a scientific method used to measure the heat transferred in a physical or chemical process, such as a chemical reaction, phase change, or solution formation.
02

Types of Calorimeters

Two commonly used calorimeters include: 1. Coffee Cup Calorimeter: A simple, constant pressure calorimeter often used in educational environments. It's typically comprised of two styrofoam cups for insulation. 2. Bomb Calorimeter: A more sophisticated, constant volume calorimeter. A known mass of a substance is combusted in the presence of an excess of oxygen gas. The heat produced measured is then used to calculate the heat of combustion.
03

Importance of Calorimeter's Heat Capacity

The heat capacity of a calorimeter is crucial because it allows for the accurate calculation of the amount of heat transferred in a process. The calorimeter absorbs a portion of the heat, and without knowledge of its heat capacity, we could not correct for this and would obtain inaccurate results.
04

Determining Heat Capacity of a Calorimeter

Calorimeter's heat capacity can be determined by adding a known quantity of hot water to a known quantity of cold water within the calorimeter and then applying the first law of thermodynamics to solve for calorimeter heat capacity. This is done either using a reaction with a known heat of reaction or a substance with a known specific heat capacity.

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Most popular questions from this chapter

A quantity of \(2.00 \times 10^{2} \mathrm{~mL}\) of \(0.862 \mathrm{M} \mathrm{HCl}\) is mixed with \(2.00 \times 10^{2} \mathrm{~mL}\) of \(0.431 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) in a constant-pressure calorimeter of negligible heat capacity. The initial temperature of the \(\mathrm{HCl}\) and \(\mathrm{Ba}(\mathrm{OH})_{2}\) solutions is the same at \(20.48^{\circ} \mathrm{C}\). For the process $$ \mathrm{H}^{+}(a q)+\mathrm{OH}^{-}(a q) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l) $$ the heat of neutralization is \(-56.2 \mathrm{~kJ} / \mathrm{mol}\). What is the final temperature of the mixed solution?

Portable hot packs are available for skiers and people engaged in other outdoor activities in a cold climate. The air-permeable paper packet contains a mixture of powdered iron, sodium chloride, and other components, all moistened by a little water. The exothermic reaction that produces the heat is a very common one- the rusting of iron: $$ 4 \mathrm{Fe}(s)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s) $$ When the outside plastic envelope is removed, \(\mathrm{O}_{2}\) molecules penetrate the paper, causing the reaction to begin. A typical packet contains \(250 \mathrm{~g}\) of iron to warm your hands or feet for up to \(4 \mathrm{~h}\). How much heat (in \(\mathrm{kJ}\) ) is produced by this reaction? (Hint: See Appendix 2 for \(\Delta H_{\mathrm{f}}^{\circ}\) values.

The combustion of \(0.4196 \mathrm{~g}\) of a hydrocarbon releases \(17.55 \mathrm{~kJ}\) of heat. The masses of the products are \(\mathrm{CO}_{2}=1.419 \mathrm{~g}\) and \(\mathrm{H}_{2} \mathrm{O}=0.290 \mathrm{~g} .\) (a) What is the empirical formula of the compound? (b) If the approximate molar mass of the compound is \(76 \mathrm{~g}\), calculate its standard enthalpy of formation.

The convention of arbitrarily assigning a zero enthalpy value for the most stable form of each element in the standard state at \(25^{\circ} \mathrm{C}\) is a convenient way of dealing with enthalpies of reactions. Explain why this convention cannot be applied to nuclear reactions.

A quantity of 0.020 mole of a gas initially at \(0.050 \mathrm{~L}\) and \(20^{\circ} \mathrm{C}\) undergoes a constant-temperature expansion until its volume is \(0.50 \mathrm{~L}\). Calculate the work done (in joules) by the gas if it expands (a) against a vacuum and (b) against a constant pressure of 0.20 atm. (c) If the gas in (b) is allowed to expand unchecked until its pressure is equal to the external pressure, what would its final volume be before it stopped expanding, and what would be the work done?
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