Temperature and Heat Study Guide (page 2)
In this lesson, we will shift our attention from forces to energy, and we will learn about heat as a form of energy, characterized by temperature We will also discuss the various types of heat transfer between objects, and tha effect of heat transfer on solids, that is, thermal expansion.
Temperature and Temperature Scales
Compared to simpler quantities such as distance and mass, temperature is more difficult to explain because it is more abstract. It's not as visible or tangible a measurement as is mass. All things constant, the larger the object, the more the mass, or volume; that is, the larger the size of the cube, the greater the volume. Temperature, on the other hand, characterizes a state of the matter, and a thermometer placed in contact with an object at different times can show different values although the object itself remains the same. Temperature measures the motion of the particles in the matter, and we will learn about this dependence later in the book. An increase in the speed of motion produces an increase in the temperature.
Historically, only a few temperature scales have been defined and adopted, the most common ones being the Celsius, or centigrade scale, and the Fahrenheit scale. Both of these scales depend on some reference points. The Celsius scale defines 0° C as the temperature where water freezes at normal atmospheric pressure, and defines 100°C as the boiling point of water. One one-hundredth of this interval is called a Celsius degree. As you can see, the scale was defined taking into consideration a specific substance—water and its properties at two standard points (0 and 100).
The Fahrenheit scale considers the same specific substance, but the standard points were assigned at the values of 32 (freezing) and 212 (boiling). So, not only does the Fahrenheit scale have an offset with respect to Celsius, it's not a centigrade (or l00-grade) scale (212 – 32 = 180). Therefore, the conversion between the two is not through a conversion factor but is a linear dependence.
If we call the Celsius temperature t(°C) and the Fahrenheit temperature t(°F), then, as we said, there is an offset between the two and a difference in the degrees. The Celsius degree is larger than the Fahrenheit degree, for example, 100 degrees Celsius equals 180 degrees Fahrenheit.
In order to eliminate the dependence on the specific substance and standard points, a more universal scale is needed. The result of that search is the absolute Kelvin scale, and its definition is based on empirical observation of the temperature dependence of the pressure of a gas at a constant volume. By lowering the temperature of a gas, the pressure is shown to decrease linearly. The pressure cannot be measured for very low temperatures where the gas liquefies, but if we extrapolate the data, we can obtain the point where the pressure becomes zero.
This point is called the zero absolute temperature, and the scale defined is Kelvin temperature scale. We can think about the Kelvin scale as an offset of the Celsius scale by –273.15 degrees; hence, the conversion between the two scales is stated simply as:
T(K) = t(°C) + 273.15
When we talk about the state of an object, we say the temperature is so many degrees Celsius (°C) or Fahrenheit (°F). When we talk about a process, a process whereby temperature increases or decreases by a value, then one says the temperature changed by so many Celsius or Fahrenheit degrees.
Find the conversion from Fahrenheit to Kelvin scale. Find the value of 32° F in Kelvin degrees and check your work.
We have already determined that the conversIOn between Fahrenheit and Celsius is:
We can use this equation to solve for t(°C) as a function of t(°F) and then use the result in T(K).
With this expression, we go back to T(K) as a function of t(°C).
t(°F) = 32°F
T(K) = 273.15K
t(°C) = 0°C
The result agrees with our initial discussion about the Kelvin scale and its direct connection with the Celsius scale.
For a system to change its temperature, the object exchanges energy with its surroundings. We call this energy heat, and we measure heat in joules (from James Joule, 1818–1889).
We have seen previously that temperature is specific to a certain state. Therefore, we call temperature a quantity of state. In contrast, heat is an energy flow established when there is a thermal contact between two different temperature states. We call heat a quantity of process.
Heat is related to a measure of the change in the motion and interaction of the particles, which is called internal energy. Internal energy is equal to the kinetic energy and the potential energy of the particles. Kinetic energy measures the translational, rotational, and vibrational motion of the atoms and molecules in the object, whereas potential energy measures the interaction between the particles. When the temperature increases, the kinetic energy increases. One way to accomplish this process is by exchanging heat with the object.
Heat will flow freely from a high-temperature object to a low-temperature object because of the difference in temperature.
We have seen why heat flows; let's learn now about the means of this flow. On an atomic scale, materials in different states are built differently: Solids have an internal structure, and the atoms and molecules are bound through strong bonds. Disturbing the bonds in one place will create a disturbance in the lattice, and so we end up with a propagation of the initial effect. Liquids and gases are different in the sense that the structure is not isotropic (the same in all directions, or completely absent as in the case of gases). Hence, a disturbance in one side of the container with fluid will spread differently than it does in a solid.
There are three types of heat transfer: conduction, convection, and radiation.
In the case of conduction, the heat is transferred through the material itself, and a difference in temperature between different sides of the object is required. On an atomic scale. the particles on the side of the object where the temperature is larger are characterized by a greater kinetic energy. While moving, they will collide with slower particles and, in the process, lose some of their kinetic energy to the slower moving particles that now accelerate.
Metals are good thermal conductors because of the free electrons that move through the lattice. Other materials such as glass, plastic, and paper are poor conductors due to the light interaction between constituent particles. Still other materials, such as gases, are isolators due to the large distance between particles.
The situation is different in convection, where the transfer of heat happens due to movement of the substance through space. Consider forcing warm air into a room through floor-level inlets. What happens in time? The warm air rises, and the cold air sinks. Warm air has atoms and molecules that move faster, and they are farther apart; therefore, the density is less than the density of cold air. And, as we have seen in the last lesson, a lower density material will be buoyed up.
A different process that does not require contact is called radiation. With this transfer, heating is accomplished by electromagnetic radiation. Every object with a temperature more than zero absolute Kelvin radiates infrared radiation, which in turn is absorbed by other objects and increases their temperature.
Heat and Temperature Change
An object that is warmed up will experience an increase in temperature. The variation of temperature is different depending on the nature of the object. To characterize this dependence, we define two coefficients: the heat capacity and the specific heat. The coefficients of heat capacity and specific heat for different materials are tabulated.
If a quantity of heat Q is transferred to a substance thereby increasing its temperature by ΔT = T2–T1 the heat capacity is defined to be:
This coefficient is not dependent on the mass of the specific object and is measured in J/K or J/C°
If a quantity of heat Q is transferred to a substance of mass m and is increasing its temperature by ΔT = T2 – T1 , the specific heat is defined to be:
We can also rewrite the equation as:
Q = m · c · (T2 – T1)
Using this equation, we define a new unit for heat called the calorie. One calorie (1 cal) is the heat necessary to be transferred to 1 g of water to increase its temperature by one Celsius degree (from 14.5 to 15.5°C). The conversion from calories to joules is:
1 cal = 4.186 J
The nutritional calorie that you see on food labels is actually 1,000 calories and is symbolized by C.
1 C = 1,000 cal
Another usual unit is one British thermal unit (BTU), which is the heat necessary to be transferred to 1 pound of water to raise its temperature from 63 to 64° F.
Some containers are built such that they make good isolators. The heat transferred to a fluid is completely exchanged with the fluid and none lost to its surroundings. Such containers are called calorimeters, and the study of the heat exchange in these systems is called calorimetry. In this case, the heat coming from a hot reservoir is completely transferred to the cold reservoir:
Qhot = –Qcold
The minus sign indicates that the system that cools down looses energy (heat is coming out), and therefore, the heat transfer is negative. This is called the calorimetric equation.
Some of the most usual materials encountered and their specific heats are shown in Table 11.1.
An aluminum piece of 400 g is placed in a container that holds 100 g water at 80° C. The water cools down to 20° C. In the process, the aluminum piece gets warmer and reaches a temperature of 45° C. What was the initial temperature of the aluminum piece? Consider the only heat exchange to be between the aluminum and the water.
First, convert all quantities to 51 units. Next, set the equations and solve for the initial temperature.
- mal = 400 g = 400 g ·1 kg/1,000 g = 0.4 kg
- mwater = 100 g = 100 g · 1 kg/1,000 g = 0.1 kg
- twater initial = 80° C
- twater final = 20° C
- tAl final = 45° C
- tal initial = ?
Because the system is thermally isolated, the heat released by the water is absorbed by the aluminum.
- Qwater = – Qaluminum
- Qwater = mwater· Cwater · (twater final – twater initial)
- Qwater = 0.1 kg · 4,186 J/kg · °C · (80° C – 20° C)
- Qwater = 0.1· 4,186 J · 60
- Qwater = 2.5 . 104 J
- –Qaluminum = –mAl · cAl · (tAl final – tAl initial)
- –Qaluminum == – 0.4 kg . 900 j/kg °C · (45 – tAl initial)
- 2.5 · 104 J = 0.4 · 900 J/ °C · (45 – tAl initial)
- 2.5 · 104 J/(0.4 · 900 J/°C) = (45 – tAl initial)
- 2.5 · 104°C/(0.4 · 900) = (45 – tAl initial)
- 70 °C = (45 – tAl initial)
- tAl initial = (45 – 70°)C = – 25°C
- tAl initial = – 25°C
Heat and Phase Change
Is the heat absorbed or released by an object always changing its temperature? The answer is no. Think, for example, about boiling water: Once the water boils, there is another phenomenon taking place called vaporization. So, in this example, the heat absorbed is used first to increase the water temperature and then to change the phase from liquid to gas. This transformation is called a phase change, and in our example, the pressure is considered to be constant because the transformation takes place in open space. The heat exchanged when the liquid is experiencing a phase change is called latent heat. Measurements show that in a phase change, the temperature is constant until the transformation to another phase happens in the bulk of the substance.
There are a few processes regarding phase changes: from solid to liquid and the inverse. These are called melting and freezing. Together, they are called fusion; specifically, they are vaporization at evaporation or condensation, and sublimation at the phase change from solid to gas or gas to solid.
The heat exchanged in a phase change has the same value regardless of the direction of change: The coefficient of latent heat of melting is equal to the coefficient of latent heat of freezing.
In many books, these processes are summarized in a graph (see Figure 11.2) of the temperature dependence of the absorbed energy (heat).
If you interpret the graph, you will see that the slopes of the graphs for ice and for steam are almost the same. All things being equal, this translates to the ratio of heat and temperature change, which gives us the specific heat. If you check with the table, you will see the two coefficients close in value. How about the water phase? The slope is smoother in this case, as shown also by the numbers in the coefficients of specific heat.
In a calorimetry measurement, the hot reservoir or/and the cold reservoir heat might contain terms similar to m · c · ΔT as we have already seen before. However, there are new terms involving a process with no temperature change, a process called phase transition. Let's try to define a measure of the transition from the point of view of heat transfer.
Coefficient of Latent Heat
If a quantity of heat Q is transferred to a substance of mass m and the substance has a phase change, then the coefficient of latent heat L is given as follows:.
This coefficient depends on the nature of the substance and on the type of phase change occurring.
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