Heat, Energy, and Work
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Figure 1. Useful Vocabulary
In chemistry there have been ideas of heat being referred to as a “flow” which introduced ways for measuring the amount of energy transferred as heat. This branch of science is called calorimetry. One way of measuring the amount of energy transferred as heat is with the use of an ice calorimeter, which includes a bath containing water and ice. When heat is added to the bath from the system some of the ice melts resulting in a higher volume in water. As a result, the amount of heat transferred is directly proportional to the volume of contents of the caliorimeter. Heat is related to the specific heat, mass, and temperature of an object. Specific heat of an object is “the amount of heat required to increase the temperature of a 1-g mass by 1C,”2 and can be written as q=McsT. Before the middle of the 19th century work, heat and energy all had the same unit, the joule and now we also use the calorie (1cal= 4.184J), another unit of energy.
It's important to study the differences of “heat”, “energy”, “work”, and “temperature” because we often used them in our everyday language in different way than how they are used in a scientific manner. One example of a common misconception incorporates the inter-changable uses of work and heat. Many students say that they are “heating up” their food in a microwave however this “heating up” process involves minimal heat. The energy in the waves does work on the water molecules in the food causing them to bump into each other which creates small amounts of "heat" (the average kinetic energy increases resulting in an increase in temperature). In contrast, when boiling a pot of water, the heat from the hot burner is conducted through the base of the pot causing uniform convectional currents throughout the water.
Heat and Temperature: Heat or thermal energy is a process by which energy is exchanged with the surroundings of a system3, in particular a diathermic system (a system that allows thermal energy exchange). In other words, heat causes changes in the internal energy of a system. Heat is what can increase a system’s internal energy, where internal energy is the sum of all kinetic and potential energies in a system, without mechanical interaction2. In mathematical terms, U=KEsystem+PEsystem where KEsystem is the energy of motion in the system and PEsystem is the position dependent, stored energy of a system.
While heat is defined as the spontaneous flow of energy from one object to another as a result of temperature differences; temperature itself is a measure of how easily an object gives up its energy to its surroundings spontaneously3.Heat prefers to go from a system with high temperature to a system with low temperature until the two systems reach a thermal equilibrium. This thermal equilibrium is where both systems reach the same temperature and there is no more heat flowing between them, a condition appropriately named- zero heat flow.Why does heat flow from high to low temperature? The answer lies in the the way that heating functions as a form of energy transfer through thermal motion, or disorderly molecular motion1. All molecules show molecular motion as seen in this video (with some exceptions for ideal gases or perfect crystals at 0 K). For a system with high temperature, molecules move rapidly in random motion. Alternately, for a system with low temperature, molecules move slowly also in random motion. This is where thermal energy comes in. As we have previously learned, Kinetic energy is proportional to velocity squared. In a set of two systems, the system with the higher temperature has more rapid moving molecules, and therefore, a greater kinetic energy. Similarly, the system with the lower temperature has slower molecules, with a lower Kinetic Energy. Therefore, the molecules in the higher temperature system stimulates the molecules the lower temperature system to move faster, thereby, increasing the energy of the lower temperature system via heat transfer1.
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Introduction:
Thermodynamics is the branch of chemistry that gives us guiding principles on energy transfers among all matter. There are two ways to change the energy of a system and that is with work or heat. In other words,work and heat are two categories of energy transfers that result in changes in a system's internal energy3. Work is done when energy is changed through mechanical contact; in contrast, heat is the mechanism of transfer when energy is changed through thermal contact. With ‘U’ representing the internal energy, ‘q’ representing heat, and ‘w’ representing work, the first law of thermodynamics can be written as U=q+w. Although this equation looks relatively simple, many Chemistry 260 students confuse these terms for one another due to similar terminology. This exploration will further explain the differences between energy, work, and heat and how these terms are related to each other.
Heat and Work vs. Energy:
As briefly mentioned in the introduction, heat and work are different from the concept of energy because they are processes of changing a system’s internal energy. Heat and work are seen as different processes because they have different ways of exchanging energy. Remember, work is done when energy is changed through mechanical contact whereas heat is transferred when energy is changed through thermal contact.
Work is defined as “the product of the external force on a body times the distance through which the force acts.”2 Work can also be written asW=Fd where ‘d’ is the distance the object has traveled and ‘F’ is the amount of force applied in the direction of movement.2 Algebraic evidence of how work can change the energy of system can be seen below (Figure 2). Pressure-volume work is an important kind of mechanical work in chemistry that results “when a system is compressed or expanded under the influence of an outside pressure.”2 Pressure-volume work can be written as W= - PV and was derived from the following;
W=Fd
- We know that in gasses pressure is P=FA where ‘F’ is the force exerted and ‘A’ is the area.
- By rearranging the equation we can derive force to be F=PA.
- We also know that the change in height of a cylinder, h, equals ‘d’.
- If we substitute what we found then we end up with, W=-(P)(A)(h).
- However, since the force (pressure) opposes the direction of movement, there is negative work (loss of energy from the system) which results in, W=-(P)(A)(h)
- “The product (A)(h) is the volume change of the system, V, so the work is”2
- W= - PV
 Figure 2. A table algebraically showing how work can change the energy (kinetic or potential) of a system.
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Common Misconceptions: Enthalpy
Many students confuse the difference between heat and enthalpy. As mentioned previously heat is “the amount of energy transferred between two objects initially at different temperatures”2. Enthalpy is a defined thermodynamic state function that is strongly related to heat and was given the symbol H much like what is commonly used to imply heat. Enthalpy can be written as
where ‘U’ represents the internal energy of the system ‘P’ represents pressure and ‘V’ represents the volume of the system. When pressure is constant, a change in enthalpy is equal to heat, which can be written as 
. The change in enthalpy, also known as the “heat of the reaction”2 is an important concept because it tells us the amount of energy that is being released and absorbed in a reaction. Overall it is key to understand that 
only when pressure is constant. If pressure is not constant then the equation; 
applies.
Common Misconceptions: Heat Capacity
Another related concept and potentially confusing concept is heat capacity which is the amount of heat supplied over the temperature change, or mathematically, 
. This heat capacity is a constant and comes in many forms. Two commonly confused forms of heat capacity are Cp and Cv. The subscripts indicate what property is constant. For Cv, volume is constant and for Cp is constant. For an ideal gas, these constants relate through the equation , Cp = Cv + R where R is the gas constant. Furthermore, these heat capacity constants vary based on atom vs molecule and molecule linearity as well as based on the measured properties. Given these variations, it is important to note that each form of heat capacity is a different constant. Also, remember that heat capacity is another property of an atom or molecule, and something that can be used to calculate change in internal energy or heat.
Concept Questions:
1.The internal energy of a system can change because:
A)The surroundings can do work on the system
B) Heat can flow out of the system
C) The system can do work on the surroundings
D) Heat can flow into the system
E) All of the above
2. If a diathermic system is exothermic which statement best describes heat transfer?
A) Heat is released
B)Energy is transferred from system into surroundings
C) Energy flows into system from surroundings
D)Both A and B
E) No heat transfer is possible in a diathermic system
3. If a reaction is carried out at constant pressure, which of the following statements is correct?
A) The reaction is likely to be exothermic.
B) The heat change is equal to the enthalpy change.
C) The reaction is likely to be endothermic.
D) The heat change is equal to the change in temperature.
4. Consider the following thermodynamic properties.
i)work done on a system
ii) Heat absorbed
iii)Entropy
iv)Enthalpy
Which of these properties are state functions?
A) iii) and iv) only
B) i) only
C) i) and iii) only
D) i) and ii) only
5. What does temperature measure?
A) Amount of heat leaving the system
B) How hot or cold the system is
C) How hot or cold the system feels
D) The average kinetic energy of molecules in the system
Citations:
[1] Atkins, P. W., and Julio Paula. "Chapter 2: The First Law." Atkins' Physical Chemistry. 8th ed. Oxford: Oxford UP, 2006. 28-32. Print.
[2] Oxtoby, David W., H. P. Gillis, and Alan Campion. Principles of Modern Chemistry. 7th ed.Cengage Learning, 2011. Print.
[3] Schroeder, David V. "Heat and Work". An Introduction to Thermal Physics. Addison Wesley Longman. 2000. 17-19. Print.
Answers: (1)E, (2)D, (3)B, (4)A (5) D
C= qT
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