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    which factors affect heat transfer between a warm and a cool substance? the time it took to heat the substance, the area of contact, and the specific heats of the substances the amount of time of contact, the area of contact, and the specific heats of the substances the time it took to heat the substance, the area of the substances, and the specific heats of the substances the amount of time of contact, the area of the substances, and the specific heats of the substances

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    get which factors affect heat transfer between a warm and a cool substance? the time it took to heat the substance, the area of contact, and the specific heats of the substances the amount of time of contact, the area of contact, and the specific heats of the substances the time it took to heat the substance, the area of the substances, and the specific heats of the substances the amount of time of contact, the area of the substances, and the specific heats of the substances from EN Bilgi.

    Heat Transfer, Specific Heat, and Calorimetry – University Physics Volume 2

    5 HEAT TRANSFER, SPECIFIC HEAT, AND CALORIMETRY

    Learning Objectives

    By the end of this section, you will be able to:

    Explain phenomena involving heat as a form of energy transfer

    Solve problems involving heat transfer

    We have seen in previous chapters that energy is one of the fundamental concepts of physics. Heat is a type of energy transfer that is caused by a temperature difference, and it can change the temperature of an object. As we learned earlier in this chapter, heat transfer is the movement of energy from one place or material to another as a result of a difference in temperature. Heat transfer is fundamental to such everyday activities as home heating and cooking, as well as many industrial processes. It also forms a basis for the topics in the remainder of this chapter.

    We also introduce the concept of internal energy, which can be increased or decreased by heat transfer. We discuss another way to change the internal energy of a system, namely doing work on it. Thus, we are beginning the study of the relationship of heat and work, which is the basis of engines and refrigerators and the central topic (and origin of the name) of thermodynamics.

    Internal Energy and Heat

    A thermal system has internal energy (also called thermal energy), which is the sum of the mechanical energies of its molecules. A system’s internal energy is proportional to its temperature. As we saw earlier in this chapter, if two objects at different temperatures are brought into contact with each other, energy is transferred from the hotter to the colder object until the bodies reach thermal equilibrium (that is, they are at the same temperature). No work is done by either object because no force acts through a distance (as we discussed in Work and Kinetic Energy). These observations reveal that heat is energy transferred spontaneously due to a temperature difference. (Figure) shows an example of heat transfer.

    (a) Here, the soft drink has a higher temperature than the ice, so they are not in thermal equilibrium. (b) When the soft drink and ice are allowed to interact, heat is transferred from the drink to the ice due to the difference in temperatures until they reach the same temperature, , achieving equilibrium. In fact, since the soft drink and ice are both in contact with the surrounding air and the bench, the ultimate equilibrium temperature will be the same as that of the surroundings.

    The meaning of “heat” in physics is different from its ordinary meaning. For example, in conversation, we may say “the heat was unbearable,” but in physics, we would say that the temperature was high. Heat is a form of energy flow, whereas temperature is not. Incidentally, humans are sensitive to heat flow rather than to temperature.

    Since heat is a form of energy, its SI unit is the joule (J). Another common unit of energy often used for heat is the calorie (cal), defined as the energy needed to change the temperature of 1.00 g of water by —specifically, between and , since there is a slight temperature dependence. Also commonly used is the kilocalorie (kcal), which is the energy needed to change the temperature of 1.00 kg of water by . Since mass is most often specified in kilograms, the kilocalorie is convenient. Confusingly, food calories (sometimes called “big calories,” abbreviated Cal) are actually kilocalories, a fact not easily determined from package labeling.

    Mechanical Equivalent of Heat

    It is also possible to change the temperature of a substance by doing work, which transfers energy into or out of a system. This realization helped establish that heat is a form of energy. James Prescott Joule (1818–1889) performed many experiments to establish the mechanical equivalent of heat—the work needed to produce the same effects as heat transfer. In the units used for these two quantities, the value for this equivalence is

    We consider this equation to represent the conversion between two units of energy. (Other numbers that you may see refer to calories defined for temperature ranges other than to .)

    (Figure) shows one of Joule’s most famous experimental setups for demonstrating that work and heat can produce the same effects and measuring the mechanical equivalent of heat. It helped establish the principle of conservation of energy. Gravitational potential energy (U) was converted into kinetic energy (K), and then randomized by viscosity and turbulence into increased average kinetic energy of atoms and molecules in the system, producing a temperature increase. Joule’s contributions to thermodynamics were so significant that the SI unit of energy was named after him.

    Joule’s experiment established the equivalence of heat and work. As the masses descended, they caused the paddles to do work, , on the water. The result was a temperature increase, , measured by the thermometer. Joule found that was proportional to W and thus determined the mechanical equivalent of heat.

    Increasing internal energy by heat transfer gives the same result as increasing it by doing work. Therefore, although a system has a well-defined internal energy, we cannot say that it has a certain “heat content” or “work content.” A well-defined quantity that depends only on the current state of the system, rather than on the history of that system, is known as a state variable. Temperature and internal energy are state variables. To sum up this paragraph, heat and work are not state variables.

    Source : opentextbc.ca

    Heat and temperature (article)

    What heat means in thermodynamics, and how we can calculate heat using the heat capacity.

    Internal energy

    Heat and temperature

    What heat means in thermodynamics, and how we can calculate heat using the heat capacity.

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    Key points

    Heat,

    \text q q

    start text, q, end text

    , is thermal energy transferred from a hotter system to a cooler system that are in contact.

    Temperature is a measure of the average kinetic energy of the atoms or molecules in the system.

    The zeroth law of thermodynamics says that no heat is transferred between two objects in thermal equilibrium; therefore, they are the same temperature.

    We can calculate the heat released or absorbed using the specific heat capacity

    \text C C

    start text, C, end text

    , the mass of the substance

    \text m m

    start text, m, end text

    , and the change in temperature

    \Delta \text T ΔT

    delta, start text, T, end text

    in the equation:

    \text q = \text {m} \times \text C \times \Delta \text T

    q=m×C×ΔT

    start text, q, end text, equals, start text, m, end text, times, start text, C, end text, times, delta, start text, T, end text

    Heat in thermodynamics

    What contains more heat, a cup of coffee or a glass of iced tea? In chemistry class, that would be a trick question (sorry!). In thermodynamics, heat has a very specific meaning that is different from how we might use the word in everyday speech. Scientists define heat as thermal energy transferred between two systems at different temperatures that come in contact. Heat is written with the symbol q or Q, and it has units of Joules (

    \text J J

    start text, J, end text

    ).

    Three melting ice cubes in a puddle of water on a mirrored surface.

    Heat is transferred from the surroundings to the ice, causing the phase change from ice to water. Photo of ice cubes from flickr, CC BY 2.0.

    Heat is sometimes called a process quantity, because it is defined in the context of a process by which energy can be transferred. We don't talk about a cup of coffee containing heat, but we can talk about the heat transferred from the cup of hot coffee to your hand. Heat is also an extensive property, so the change in temperature resulting from heat transferred to a system depends on how many molecules are in the system.

    Relationship between heat and temperature

    Heat and temperature are two different but closely related concepts. Note that they have different units: temperature typically has units of degrees Celsius (

    ^\circ\text C ∘ C

    degrees, start text, C, end text

    ) or Kelvin ( \text K K

    start text, K, end text

    ), and heat has units of energy, Joules (

    \text J J

    start text, J, end text

    ). Temperature is a measure of the average kinetic energy of the atoms or molecules in the system. The water molecules in a cup of hot coffee have a higher average kinetic energy than the water molecules in a cup of iced tea, which also means they are moving at a higher velocity. Temperature is also an intensive property, which means that the temperature doesn't change no matter how much of a substance you have (as long as it is all at the same temperature!). This is why chemists can use the melting point to help identify a pure substance

    - − minus

    the temperature at which it melts is a property of the substance with no dependence on the mass of a sample.

    On an atomic level, the molecules in each object are constantly in motion and colliding with each other. Every time molecules collide, kinetic energy can be transferred. When the two systems are in contact, heat will be transferred through molecular collisions from the hotter system to the cooler system. The thermal energy will flow in that direction until the two objects are at the same temperature. When the two systems in contact are at the same temperature, we say they are in thermal equilibrium.

    Zeroth law of thermodynamics: Defining thermal equilibrium

    The zeroth law of thermodynamics defines thermal equilibrium within an isolated system. The zeroth law says when two objects at thermal equilibrium are in contact, there is no net heat transfer between the objects; therefore, they are the same temperature. Another way to state the zeroth law is to say that if two objects are both separately in thermal equilibrium with a third object, then they are in thermal equilibrium with each other.

    The zeroth law allows us to measure the temperature of objects. Any time we use a thermometer, we are using the zeroth law of thermodynamics. Let's say we are measuring the temperature of a water bath. In order to make sure the reading is accurate, we usually want to wait for the temperature reading to stay constant. We are waiting for the thermometer and the water to reach thermal equilibrium! At thermal equilibrium, the temperature of the thermometer bulb and the water bath will be the same, and there should be no net heat transfer from one object to the other (assuming no other loss of heat to the surroundings).

    Heat capacity: Converting between heat and change in temperature

    How can we measure heat? Here are some things we know about heat so far:

    Source : www.khanacademy.org

    12.3: Heat Capacity, Enthalpy, and Calorimetry

    Hess's law is that the overall enthalpy change for a series of reactions is the sum of the enthalpy changes for the individual reactions. For a chemical reaction, the enthalpy of reaction (ΔH) …

    12.3: Heat Capacity, Enthalpy, and Calorimetry

    Last updated Jan 7, 2022

    12.2: The First Law of Thermodynamics - Internal Energy, Work, and Heat

    12.4: Illustrations of the First Law of Thermodynamics in Ideal Gas Processes

    Learning Objectives

    Explain the technique of calorimetry

    Calculate and interpret heat and related properties using typical calorimetry data

    To use calorimetric data to calculate enthalpy changes.

    Heat Capacity

    We now introduce two concepts useful in describing heat flow and temperature change. The heat capacity (

    C C

    ) of a body of matter is the quantity of heat (

    q q

    ) it absorbs or releases when it experiences a temperature change (

    ΔT ΔT

    ) of 1 degree Celsius (or equivalently, 1 kelvin)

    C= q ΔT (12.3.1) (12.3.1)C=qΔT

    Heat capacity is determined by both the type and amount of substance that absorbs or releases heat. It is therefore an extensive property—its value is proportional to the amount of the substance.

    For example, consider the heat capacities of two cast iron frying pans. The heat capacity of the large pan is five times greater than that of the small pan because, although both are made of the same material, the mass of the large pan is five times greater than the mass of the small pan. More mass means more atoms are present in the larger pan, so it takes more energy to make all of those atoms vibrate faster. The heat capacity of the small cast iron frying pan is found by observing that it takes 18,140 J of energy to raise the temperature of the pan by 50.0 °C

    C small pan = 18,140J 50.0°C =363J/°C

    Csmall pan=18,140J50.0°C=363J/°C

    The larger cast iron frying pan, while made of the same substance, requires 90,700 J of energy to raise its temperature by 50.0 °C. The larger pan has a (proportionally) larger heat capacity because the larger amount of material requires a (proportionally) larger amount of energy to yield the same temperature change:

    C large pan = 90,700J 50.0°C =1814J/°C

    Clarge pan=90,700J50.0°C=1814J/°C

    The specific heat capacity (

    c c

    ) of a substance, commonly called its specific heat, is the quantity of heat required to raise the temperature of 1 gram of a substance by 1 degree Celsius (or 1 kelvin):

    c= q mΔT (12.3.2) (12.3.2)c=qmΔT

    Specific heat capacity depends only on the kind of substance absorbing or releasing heat. It is an intensive property—the type, but not the amount, of the substance is all that matters. For example, the small cast iron frying pan has a mass of 808 g. The specific heat of iron (the material used to make the pan) is therefore:

    c iron = 18,140J (808g)(50.0°C) =0.449J/g°C

    ciron=18,140J(808g)(50.0°C)=0.449J/g°C

    The large frying pan has a mass of 4040 g. Using the data for this pan, we can also calculate the specific heat of iron:

    c iron = 90,700J (4,040g)(50.0°C) =0.449J/g°C

    ciron=90,700J(4,040g)(50.0°C)=0.449J/g°C

    Although the large pan is more massive than the small pan, since both are made of the same material, they both yield the same value for specific heat (for the material of construction, iron). Note that specific heat is measured in units of energy per temperature per mass and is an intensive property, being derived from a ratio of two extensive properties (heat and mass). The molar heat capacity, also an intensive property, is the heat capacity per mole of a particular substance and has units of J/mol °C (Figure

    12.3.1 12.3.1 ).

    Figure 12.3.1 12.3.1

    : Due to its larger mass, a large frying pan has a larger heat capacity than a small frying pan. Because they are made of the same material, both frying pans have the same specific heat. (CC BY; Mark Blaser via OpenStax).

    The heat capacity of an object depends on both its mass and its composition. For example, doubling the mass of an object doubles its heat capacity. Consequently, the amount of substance must be indicated when the heat capacity of the substance is reported. The molar heat capacity (Cp) is the amount of energy needed to increase the temperature of 1 mol of a substance by 1°C; the units of Cp are thus J/(mol•°C).The subscript p indicates that the value was measured at constant pressure. The specific heat (

    c s cs

    ) is the amount of energy needed to increase the temperature of 1 g of a substance by 1°C; its units are thus J/(g•°C).

    We can relate the quantity of a substance, the amount of heat transferred, its heat capacity, and the temperature change either via moles (Equation

    12.3.3 12.3.3 ) or mass (Equation 12.3.4 12.3.4 ): q=n c p ΔT (12.3.3) (12.3.3)q=ncpΔT where n n

    is the number of moles of substance and

    c p cp

    is the molar heat capacity (i.e., heat capacity per mole of substance), and

    ΔT= T final

    Source : chem.libretexts.org

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