Heat and First Law of Thermodynamics Vocabulary calorie kilocalorie Calorie mechanical equivalent of heat heat thermal energy internal energy specific heat calorimetry calorimeter conservation of energy water equivalent bomb calorimeter change of phases heat of fusion heat of vaporization latent heat evaporation system closed system open system isolated first law of thermodynamics isothermal heat reservoir quasistatically adiabatic isobaric isochoric free expansion molar specific heats degrees of freesom principle of equipartion of energy conduction thermal conductivity conductors insulators convection radiation emissivity Stefan Boltzmann equation Stefan Boltzman constant solar constant thermography 19-1 Heat as Energy Transfer • Heat spontaneous flows from the hotter object to the cooler object •This will continue until equilibrium is reached •Thermometers work this way •heat- is a term that is used inconsistently in real life •We generally speak of the flow of heat- heat flows from objects of higher temperatures to lower temperatures •in the 18th century the model of heat pictured heat flow as a movement of a fluid substance called caloric •caloric fluid could never be detected •19th century- that the various phenomena associated with the heat could be described consistently without the need to use the fluid model •New model viewed heat as being akin to work and energy •common unit for heat still used today is named after the caloric •calorie- cal- the amount of heat necessary to raise the temperature of 1 gram of water by 1˚C •kilocalorie- kcal- 1000 calories- 1 kcal is the heat is needed to raise 1 kg of water by 1˚C •Calorie- C- this is the unit that is used for food- dietary calorie •The idea that heat is related to energy was pursued by various scientist in the 1800’s- James Prescott Joule (1818-1889) •He performed experiments that aided our understanding of heat, work, transfer of heat •www.youtube.com/watch?v=5yOhSIAIPRE Name:_____________________________________ Num:________ 1 RoessBoss
Transcript
Heat and First Law of Thermodynamics
Vocabularycalorie kilocalorie Calorie mechanical equivalent of
heatheat thermal energy internal energy specific heat
calorimetry calorimeter conservation of energy water equivalent
bomb calorimeter change of phases heat of fusion heat of vaporization
latent heat evaporation system closed system
open system isolated first law of thermodynamics
isothermal
heat reservoir quasistatically adiabatic isobaric
isochoric free expansion molar specific heats degrees of freesom
Stefan Boltzman constant solar constant thermography
19-1 Heat as Energy Transfer• Heat spontaneous flows from the hotter object to the cooler object•This will continue until equilibrium is reached•Thermometers work this way•heat- is a term that is used inconsistently in real life•We generally speak of the flow of heat- heat flows from objects of higher temperatures to lower temperatures•in the 18th century the model of heat pictured heat flow as a movement of a fluid substance called caloric•caloric fluid could never be detected•19th century- that the various phenomena associated with the heat could be described consistently without the need to use the fluid model•New model viewed heat as being akin to work and energy•common unit for heat still used today is named after the caloric
•calorie- cal- the amount of heat necessary to raise the temperature of 1 gram of water by 1˚C•kilocalorie- kcal- 1000 calories- 1 kcal is the heat is needed to raise 1 kg of water by 1˚C•Calorie- C- this is the unit that is used for food- dietary calorie
•The idea that heat is related to energy was pursued by various scientist in the 1800’s- James Prescott Joule (1818-1889)•He performed experiments that aided our understanding of heat, work, transfer of heat•www.youtube.com/watch?v=5yOhSIAIPRE
• The falling weight causes the paddle wheel turn and do work on the water•The friction between the water and the paddle wheel causes the temperature of the water to rise slightly•The same temperature rise could also be obtained by heating the water on a hot stove•Joule determined that a given amount of work done was always equivalent to a particular amount of heat input•Quantitively - 4.186 J- of work was found to be equivalent to 1 calorie•This relationship is known as the mechanical equivalent
of heat
Equation Box 19-1
•As a result of these experiments and others- scientist interpret heat not as a substance and not even as a form of energy•Heat is a transfer of energy•When energy flows from a hot object to a cooler one, it is energy that is being transferred from the hot to the cold object•heat- is energy that is transferred from one body to another because of difference in temperatures•SI units for energy are Joules•Calories, calories, and kilocalories are also used•calorie is defined in terms of the joule rather than in the properties of water•The later is still handy to remember- 1 cal raises 1 g of water by 1˚C or 1 kcal raises 1 kg of water by 1˚C
19-2 Internal Energy• The sum of all the energy of all the molecules in an object is called its thermal energy or internal energy•We can use those terms interchangeably• heat content of a body is also used for this purpose, but it is not a good term because it can be confused with heat itself•Heat is not energy a body contains but rather refers to the amount of energy transferred from one body to another at different temperatures•Distinguishing Temperature, Heat, and Internal Energy•Using the kinetic theory we can make clear distinctions between temperature, heat, and internal energy
•temperature (kelvins) is a measure of the average kinetic energy of individual molecules•Thermal energy and internal energy refer to the total energy of all the molecules in the object•Two equal mass hot ingots of iron may have the same temperature, but two of them may have twice as much thermal energy as one does•Heat- refers to the transfer of energy (such as thermal energy) from one object to another due to difference in temperature•The direction of heat flow depends on their temperatures and not on how much internal energy each has•Internal Energy of an Ideal Gas• The internal energy is U- the sum of translational kinetic energy of all the atoms•n- the number of moles of an ideal monatomic (one atom per molecule) gas•The sum is equal to the average kinetic energy per molecule times the total number of molecules N
Equation Box 19-2
• We can rewrite this as
Equation Box 19-3
• The internal energy of an ideal gas depends only on temperature and the number of moles of gas•If the gas molecule contains more than one atom, then the rotational and vibrational energy of the molecule must also be taken into account•The internal energy will be greater at a given temperature than for a monoatomic gas, but it will function only of temperatures for an ideal gas•The internal energy of real gases also depends mainly on temperature- but they deviate from ideal gas behavior- the internal energy also depends somewhat on pressure and volume•The internal energy of liquids and solids is quite complicated for it includes electrical potential energy associated with the forces (chemical bonds_ between atoms and molecules•Remember that gas molecules are spread out farther than solids and liquids so the attraction between them is less than in solids and liquids
19-3 Specific Heat• If heat is put into an object, its temperature rises•How much a temperature rises depends on quantities•in the 18th century, experimenter had reorganized that the amount of heat Q required to change the temperature of a given material is proportional to the mass of m of the material and to the the temperature change ∆T•The equation is
Equation Box 19-4
• c is the quantity characteristic of the material and is called specific heat•The units for C are J/kg•˚C or kcal/kg•˚C•For water the values are
Equation Box 19-5
•This is due to the original definition of cal and joule- it takes 1 kcal of heat to raise the temperature of 1 kg of water by 1˚C•The value of c depends to some extent on temperature •c is considered to be fairly constant on most cases•Water has one of the highest specific heats of all substances, which makes it an ideal substance for hot water space heating systems and other uses that require a minimal drop in temperature for a given amount of heat transfer
Specific Heats (1atm constant pressure and 20˚C unless otherwise stated)
Specific Heats (1atm constant pressure and 20˚C unless otherwise stated)
Glass 0.2 840
Iron or Steel 0.11 450
Lead 0.031 130
Marble 0.21 860
Mercury 0.21 860
Silver 0.056 230
Wood 0.4 1700
Water
Ice (-5˚C) 0.5 2100
Liquid (15˚C) 1 4186
Steam (110˚C) 0.48 2010
Human Body 0.83 3470
Protein 0.4 1700
19-4 Calorimetry• When different parts of isolated system are at different temperatures, but are placed in thermal contact, heat will flow from the part at higher temperature to the part at lower temperature
• If the system is completely isolated, no energy can flow into or out of it•Conservation of energy plays an important role in this concept•Heat lost= heat gained•The exchange of energy is the basis for the technique known as calorimetry, which is quantitative measurement of heat exchange•To make the measurements you use a calorimeter
•Water is generally used to insulate it so that only a minimal amount of heat is exchanged with the outside•One use of the calorimeter is to determine specific heat values
•In a “method of mixtures” technique a sample of the substance is heated to a high temperature, which is measured, and then quickly placed in the cool water of the calorimeter. Then you measure the final temperature and perform the calculations•mcalCcal- is often called the water equivalent of the calorimeter- it is numerically equal to the mass of water (in kilograms) that would absorb the same amount of heat
• A bomb calorimeter is used to measure the heat released when a substance burns•you use this to determine the energy in food to determine their energy Calories•A carefully weighed sample of the substance, together with an excess amount of oxygen at high pressure is placed in a sealed container :the bomb”. The bomb is placed in the water of the calorimeter and a fine wire passing into the bomb is then heated briefly, which cause the mixture to ignite
19-5 Latent Heat• When a material changes phase from solid to liquid to gas, a certain amount of energy is involved in this change of phase• The heat required to change 1.0 kg of a substance from the solid to a liquid is called the heat of fusion•The heat required to change a substance from liquid to vapor phase is called the heat of vaporization• Values for the heats and of fusion and vaporization which are called latent heats can be found on tables for various substances•The heats of vaporization and fusion also refer to the
amount of heat released by a substance when it changes from a gas to a liquid, or from a liquid to a solid•The heat involved in a change in phase depends not only on the latent heat, but also on the total mass of the substance
Equation Box 19-6
•L is the latent heat of the particular process and substance n is the mass of the substance and Q is the heat required or given off during the phase change•When we make use of kinetic theory to see why energy is needed to melt or vaporize a substance
•at the melting point, the latent heat of fusion does not increase the kinetic energy (and the temperature) of the molecules in the solid, but instead is used to overcome the potential energy associated with the forces between the molecules•Work must be done against these attractive forces to break the molecules loose from their relatively fixed positions in the solid so they can freely roll over one another in the liquid phase•Energy is required for molecules held close together in the liquid phase to escape into the gaseous phase•This process is more violet reorganization of the molecules than is melting (the average distance between the molecules is greatly increased) and hence the heat of vaporization is generally much greater than the heat of fusion for a given substance•The latent heat to change a liquid to a gas is needed not only at the boiling point•Water can change from liquid to gas phase even at room temperature•This process is called evaporation•The value of heat of vaporization increases slightly with a decrease in temperature•When water evaporates it cools since the energy required comes from the water itself•So the internal energy and therefore its temperature drops
19-6 The first Law of Thermodynamics• We have focused on internal energy but work is also involved in the thermodynamic process•Work is done when energy is transferred from one body to another by mechanical means•Heat is a transfer of energy from one body to a second body at a lower temperature•Heat is much like work•‘To distinguish them, heat is defined as the transfer of energy due to a difference in temperature, whereas work is a transfer of energy that is not due to a temperature difference•In thermodynamics we will refer to a particular system•A system is any object or set of objects that we wish to consider•Everything else in the universe is referred to as the environment•A closed system is one for which no mass enters or leaves (but energy may be exchanged with the environment)•Open systems mass may enter or leave (as well as energy)•Many (idealized) systems we study in physics are closed systems•many systems are open like plants, animals, etc since the exchange occurs with the environment•A closed system is said to be isolated if no energy in any form passes across its boundaries; otherwise it is not isolated•Internal energy of a system is the sum total of all the energy of the molecules of the system•Internal energy of a system would be increased if work were done on the system, or if heat were added to it
•Internal energy would be decreased if heat flowed out of the system or if work were done by the system on something else•Conservation of energy, it is reasonable to propose an important law; the change in internal energy of a closed system ∆U, will be equal to the heat added to the system minus the work done by the system
Equation Box 19-7
• This is the first law of thermodynamics•Q is the net heat added to the system•W is the net work done by the system• You must be careful with the sign convention for Q and W•Heat added is +•Heat lost is -•Work on system is -•Work by system in +•If work is done on the system, W will be negative and U will increase•If Q is positive for heat added to the system, so if heat leaves the system Q is negative•Equation 19-7 is the First law of thermodynamics•it is one of the great laws of physics•its validity rest on the experiments of joule in which no exceptions have been seen•since Q and W represent energy transfered into or out of the system, the internal energy changes accordingly•The first law of thermodynamics is a great and broad statement of the law of conservation of energy•It is worth noting that the conservation of energy law was not formulated until the 19th century, for it depended on the interpretation of heat as a transfer of energy•The equation above works for a closed system•It applies to an open system if we take into account the change in internal energy doe to the increase or decrease in the amount of matter•For an isolated system, no work is done and no heat enters or leaves the system•so W= Q= 0 and ∆U = 0•Internal energy is a property of a system; work and heat are not•a given system in a particular system is said to have a certain amount of internal energy, U
•This cannot be said of work or heat•A system in a given state does not have a certain amount of heat or work•When work is done on a system (such as compressing a gas) or when heat is added or removed from a system, the state of the system changes•Work and heat are involved in thermodynamic processes that can change the system from one state to another, they are not characteristic of the state itself, as are Pressure, Volume, Temperature, internal energy, mass, and moles•U is a state variable- a variable that depends only on the state of the system and not how the system arrived in that state
Equation Box 19-8
•U1 and U2 represent the internal energy of the system in states 1 and states 2 ad Q and W are the heat added to the system and work done by the system in going from state 1 and state 2
19-7 Applying the First Law of Thermodynamics; Calculating Work• Isothermal Processes (∆T = 0)•First consider an idealized process that is carried out at constant temperature- isothermal process•If the system is an ideal gas, then
Equation Box 19-9
•so for a fixed amount of gas kept at constant temperature, PV= constant•Process follows a curve like AB on the PV diagram which is the curve for PV= Constant•Each point on the curve such as point A, represents the state of the system as a given moment- that is pressure P and Volume V•At a lower temperature, another isothermal process
would be represented by a curve like A’B’ (the product PV= nRT= constant is less when T is less)•The curves are referred to as isotherms•assume that the gas is enclosed in a container fitted with a moveable piston
•The gas is in contact with a heat reservoir (a body whose mass is large that, ideally, its temperature does not change significantly when heat is exchanged with out system)•Also assume that the process of compression (volume decrease) or expansion (volume increase) is done quasistatically (almost statically), by which we mean extremely slowly, so that all of the gas stays in equilibrium at the same constant temperature•IF the gas is initially in a state represented by point A and the amount of heat Q is added to the system, the system will move to another point B on the diagram•If the temperature is to remain constant, the gas must expand and do an amount of work W on the environment (it exerts a force on the piston and moves it through a distance) •The temperature and mass are kept constant, the internal energy does not change
Equation Box 19-10
•Hence by the first law of thermodynamics
Equation Box 19-11
•The work done by the gas in an isothermal process equals the heat added to the gas•Adiabatic Process Q=0•An adiabatic process is one in which no heat is allowed to flow into or out of the
system Q=0•This situation can occur if the system is extremely well insulated or the process happens so quickly that heat- which flows slowly-has no time to flow in or out•The very rapid expansion of gases in an internal combustion engine is one example of a process that is very nearly adiabatic expansion of an ideal gas follows a curve like the labeled AC•Q= 0 we use the equation and find that ∆U= -W
•The internal energy decreases if the gas expands; hence the temperature decreases as well because
•An adiabatic PV curve is steeper than an isotherm- the product PV= nRT is less at point C than at point B (curve AB is for an isothermal process for which ∆U= - and ∆T = 0), work is done on the gas, and hence the internal energy increases and the temperature rises•In a diesel engine, the rapid adiabatic compression reduces the volume by a factor of 15 or more; the temperature rise is so great that the air fuel mixture ignites spontaneously•Isobaric and Isochoric
• Isothermal and adiabatic process are just two possible process that can occur•The other simple thermodynamic process are illustrated on the PV Diagram• Isobaric- is one in which the pressure is kept constant, so the process is represented by a horizontal straight line on the PV diagram•Isochoric or Isovolumeteric process is one in which the volume does not change•In theses and in all other processes the first law of
thermodynamics holds•Work Done in Volume Changes•We often want to calculate the work done in a process•Suppose we have a gas confined to a cylindrical container fitted with a moveable piston•We must always define exactly what our system is•Our system is the gas•the containers walls and the piston are parts of the environment•let us calculate the work done by the gas when it expands quasistatically so that P and T are defined for the system at all instants•The gas expands against the piston whose area is A•The gas exerts a force F=PA on the piston, where P is the pressure in the gas•The work done by the gas to move the piston an infinitestimal displacement dl is
•Since dV= A dl•If the gas were compressed so that dl pointed into the gas, the volume would decrease and dV<0•The work done by the gas in this case would be negative, which is equivalent to saying that positive work was done on the gas, not by it•The work done in taking a system from one state to another depends not only on the initial and final states but also on the type of process (or path)•The amount of heat added or removed in taking a system from one state to another depends not only on the initial and final states but also on the path or the process•Free Expansion•one type of adiabatic process is called the free expansion in which a gas is allowed to expand in volume adiabatically without doing any work•no heat flows in or out (Q=0) and no work is done because the gas does not move any other objects•Thus Q=W=0 and by the first law of thermodynamics ∆U=0•The internal energy of a gas does not change in a free expansion•for an ideal gas ∆T= 0•also since U depends only on T•Experimentally the free expansion has been used to determine if the internal energy of real gases depends only on T•The experiments are very difficult to do accurately, but it has been found that the temperature of a real gas drops very slightly in a free expansion•Thus the internal energy of real gases does depend, a little, on pressure or volume as well as on temperature•free expansion could not be plotted on a PV diagram because the process is rapid, not quasistative•immediate states are not equilibrium states, and hence pressure is not clearly defined
19-8 Molar Specific Heats for gases and Equipartition of energy• The value of the specific heat for gases depends on how the process is carried out. •Two important processes are those in which either the volume or the pressure is kept constant•Although for solids and liquids the process matters little•Molar Specific Heats for Gases•The difference in specific heats for gases is nicely explained in terms of the first law of thermodynamics and kinetic theory
•The values for specific heats can be calculated using the kinetic theory, and the results are in close agreement with experiment•molar specific heats, Cv and Cp, which are defined as the heat required to raise 1 mol of the gas by 1˚C at constant volume and at constant pressure•We simply replace the C in Q=mc∆T with the respective C that is needed•Molar Specific Heat
Equation Box 19-14
•M is the molecular mass of the gas (M= m/n)•values for molar specific heats values are nearly the same for the different gases that have the same number of atoms per molecule•Lets look at the kinetic theory and see why the specific heats of gases are higher for constant pressure processes than for constant volume processes•In both process the ∆T is the same•In the process done at constant volume, no work is done since ∆V = 0.•according to the first law of thermodynamics, the heat added (Qv) all goes into increasing the internal energy of the gas
Equation Box 19-15
•In the process carried out at constant pressure, work is done and hence the heat added Qp must not only increase the internal energy but also is used to do the work W=P∆V•Thus more heat must be added in this process than in the first process at constant volume•For the process at constant pressure we have from the first law of thermodynamics
Equation Box 19-16
•Since ∆U is the same in the two processes we can combine the two equations
•From the ideal gas Law V=nRT/P, so for a process at constant pressure we find that ∆V= nRT•Place it into the equation
Equation Box 19-18
•then cancel out like items
Equation Box 19-19
•Since the gas constant is R=8.315 J/mol•K= 1.99 cal/mol•K•our prediction we have Cp will be larger than Cv by about 1.99 cal/mol•K•Indeed this is close to what is obtained experimentally•let us now calculate the molar specific heat of a monoatomic gas using the kinetic theory model of gases•Consider a process carried out at constant volume.•Since no work is done in this process the first law of thermodynamics tells us that if heat Q is added to the gas, the internal energy of the gas changes by
Equation Box 19-20
•For an ideal monatomic gas, the internal energy U, is the total kinetic energy of all the molecules
•Since R= R=8.315 J/mol•K= 1.99 cal/mol•K, kinetic theory predicts that Cv= 2.98 cal/ mol•K for an ideal monoatomic gas•This is very close to the experimental values for monatomic gases such as helium and neon•Cp is also in agreement with the experiment•Equipartition Energy• The measured molar specific heats for more complex gases increase with the increase number of atoms per molecule•This can be explained by assuming the internal energy includes not only translational kinetic energy but other forms of energy.•A diatomic molecule has rotational along with translational kinetic energy•In molecules there is a degree of freedom which mean the number of independent ways molecules possess energy•monoatomic molecules have three degrees of freedom, since an atom can have velocity along the x, y, and z axis•A diatomic molecule has the same three degrees of freedom associated with translational kinetic energy plus two more degrees of freedom for rotational kinetic energy•19th century- scientist determined the ratio of freedoms- this is called the principle of equipartition of energy
•It states that energy is shared equally among the active degrees of freedom, and in particular each active degree of freedom of a molecule has on average an energy equal to 1/2kT•This the average energy for a molecule has on a monoatomic gas would be 3/2kT•Diatomic- 5/2kT•Hence the internal energy would be
Equation Box 19-24
•Using the previous argument we discover that the molar specific heat at constant volume would be 5/2R= 4.97 cal/mol•K in accordance with the measured values•more complex molecules have even more degrees of freedom and this greater molar specific heats•The situation was complicated- diatomics at very low temperatures have a value of only 3/2R as if it only had three degrees of freedom•At very high temperatures the Cv was 7/2 R- as if there were 7 degrees of freedom•The explanation is that at low temperatures, nearly all molecules have only translational kinetic energy•That is no energy goes into rotational energy, so only three degrees of freedom are active•At very high temperatures all five degrees of freedom are active plus two additional ones•The two new ones are associated with vibrational (pretend like the molecules are connected by a vibrating spring)•Einstein later explained the phenomenon of fewer degrees of freedom at lower temperatures through his quantum theory- energy does not tak on continuos values but is quantized- it can only have only certain values, and there is a certain minimum energy. The minimum rotational and vibrational energies are higher than for simple translational kinetic energy, so at lower temperature and lower translational kinetic energy, there is not enough energy to excite the rotational or vibrational kinetic energy•Solids•you can apply the equipartition of energy to solids as well•molar specific heat of any solid, at high temperature, is close to 3R•This is called the Dulong and Petit value after the scientist who first measured it in 1819•At high temperatures each atom apparently has six degrees of freedom although some are not active at low temperatures•each atom in a crystalline solid can vibrate about its equilibrium position as if it were connected by springs to each of its neighbors
•Thus it can have three degrees of freedom for kinetic energy and three more associated with potential energy of vibration in each of the x,y, and z directions
19-9 Adiabatic Expansion of Gas• The PV curve for the quasistative (slow) adiabatic expansion (Q=0) of an ideal gas was shown earlier•It is somewhat steeper than for an isothermal process (∆T=0), which indicates that for the same change in volume the change in pressure will be greater•The temperature of the gas must drop during an adiabatic expansion•Conversely the temperature rises during an adiabatic compression•When we use the first law and this concept we can derive a relationship. •The relation between P and V for a quasistatic adiabatic expansion or contraction is that PV is constant•This is important for heat engines.•However it is important to remember the ideal gas law PV=nRT- which continues to hold even for an adiabatic expansion (PV^¥ = constant); clearly PV is not constant- meaning T is not constant•
19-10 Heat Transfer; Conduction, Convection, Radiation• Heat is transferred from one place or body to another in three different ways; conduction, convection, and radiation•Conduction•Conduction- is a result of molecular collisions- as one end of the object is heated, the molecules move faster and faster•They collide with slower moving particles, they transfer their energy, more molecules begin to move faster•This continues until the whole body is at equilibrium•Heat conduction takes place only if there is a difference in temperature•The rate of heat flow through a substance is proportional to the difference in temperature between its ends•the rate of heat flow also depends on the size and shape of the object•let us consider the heat flow through a uniform object•It is found that the heat flow ∆Q per time interval ∆t is given by
Equation Box 19-25
•A is the cross sectional area of the object•l is the distance between the two ends
•k is the proportionality constant- thermal conductivity- which is characteristic of the material•If the material is not constant then you must use calculus to determine the values•Substances for which k is large conduct heat rapidly and are said to be good conductors•most metals fall into this category•Substances for which k is small such as fiberglass and down are poor conductors of heat and are good insulators•relative magnitudes of k can explain various phenomenon- like cooler tiles than the carpet
•For practical purposes the thermal properties of building materials are commonly specified by R values (thermal resistance)
Equation Box 19-26
•The R values of a given piece of material thus combines the thickness l and the thermal conductivity k in one number•R values increase directly with material thickness•Convection•liquids and gases are not very good conductors of heat•they can transfer heat quite rapidly by convection•convection- is the process whereby heat is transferred by the mass movement of molecules from one place to another•Conduction involves molecules (and/or electrons) moving over small distances and colliding•Convection involves the movement of molecules over large distances•Forced convection- is when air is heated and then blown by a fan into a room•Natural convection- is when nature does it itself- like hot air rising•Air expands as it is heated- density decreases- air becomes more buoyant•Warm and Cold currents of the gulf are natural convection•Radiation•Convection and conduction require the presence of matter as a medium to carry the heat from the hotter to the colder region•Radiation does not require a medium•The sun produced the radiant energy that the earth uses to sustain life•Radiation involves electromagnetic waves•For right now- it consists of visible light plus many other wavelengths that eye is not sensitive to including infrared radiation which may responsible for heating the Earth•The rate at whcih an object radiates energy has been found to be proportional to the fourth power of the Kelvin Temperature•That is a body at 2000K compared to one at 1000K radiates energy 2^4= 16 times greater•the rate of radiation is also proportional to the area A of the emitting object•So that the rate at which energy leaves the object ∆Q/∆t
• This is called the Stefan-Boltzmann equation and _____ is a universal constant called the Stefan Boltzmann constant which has the value of 5.67 E -8 W/m^2•K^4•The factor ______ is called the emissivity, is a number between 0 and 1 that is characteristic of the material•Very black surfaces have emissivity close to 1•bright shiny surfaces have emissivity close to 0- thus emit less radiation•Remember that white surfaces absorb little of the radiation- they reflect most of them unlike black surfaces•Thus a good absorber is a good emitter•Any object not only emits energy by radiation but also absorbs energy radiated by other bodies•The net flow of radiant heat from the object in relation to its environment is given by the equation
Equation Box 19-28
•A is the surface area of the object•T is the temperature•____ is the emissivity•the second temperature is that of the surroundings•The proportionality constant is the same for both emission and absorption•You can not use the above equation to determine the heating of a substance from the sun•the temperature 2 is not uniform in this case•So we treat the sun as a separate source •1350 W/m^2 is considered to be the solar constant•The atmosphere may absorb as much as 70 percent of this energy before it reaches the ground depending on cloud cover•We also have to take in account the angle at which the sun is hitting the earth.•The equation then becomes
•Thermography- uses a special instrument called a thermograph- it scans the body measuring the intensity of radiation from many points and forming a picture that resembles an X-ray•Areas where metabolic activity is high, such as tumors, can often be detected on a thermogram as a result of their higher temperature and consequent increased radiation
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