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Heat and Temperature.

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Heat and modern technology are inseparable. These glowing steel slabs, at over 1,100OC (about 2,000OF), are cut by an automatic flame torch. The slab caster converts 300 tons of molten steel into slabs in about 45 minutes. The slabs are converted to sheet steel for use in the automotive, appliance and building industries.

    • Ancient Greeks knew that matter was made up of very small particles.
    • Democritus wrote that matter was made up of tiny indivisible particles he called atoms.
      • We now know that atoms are not indivisible, but are themselves made up of even smaller particles.
      • We have identified more that 200 smaller particles that make up atoms.
    • Basic assumption is that matter is made up of tiny units of structure called atoms.
    • Atoms are neither created or destroyed during any type of chemical or physical change.
    • Arrangements of atoms determines type of entity of matter.
      • Elements are pure substances made up of only one type of atom.
      • Compounds are made up of one type of atom, but have more complex structures.
Pure substances are composed of 2 or more elements in defined proportions.
  • A molecule is the smallest particle of a compound in which all of the atoms maintain their identity.
    • maintains all of the chemical and physical properties of the compound.
    • Some atoms naturally form molecules called diatomic molecules: O2, F2, Cl2, N2, Br2,
Metal atoms appear in the micrograph of a crystal of titanium niobium oxide, magnified 7,800,000 times by an electron microscope.
Molecules Interact.
    • Cohesion.
      • Some solids and liquids attract each other and cling to each other.
      • Cohesion is when this attractive force is between like molecules.
    • Adhesion.
      • Some molecules are attracted to other molecules.
Phases of Matter.
    • Solids.
      • Defined shapes.
      • Defines volumes.
      • Molecules are fixed distances apart and have strong cohesive forces.
    • Close together.
    • Cohesive forces not as strong as in a solid.
    • Defined volume, but not a defined shape.
  • Gases
    • Weak cohesive forces.
    • High kinetic energy.
    • Molecules far apart and move in random motion
    • No fixed shape or volume.
    • Vapor is a gas that is above a liquid phase.

(A) In a solid, molecules vibrate around a fixed equilibrium position and are held in place by strong molecular forces. (B) In a liquid, molecules can rotate and roll over each other because the molecular forces are not as strong. (C) In a gas, the molecules move rapidly in random free paths.

Molecules Move.
    • All molecules have kinetic energy due to movements.
    • This kinetic energy can be in the form of:
      • Vibrational energy.
      • Rotational energy.
      • Translational energy where the entire molecule has motion.

The basic forms of kinetic energy of molecules. (A) Translational motion is the motion of a molecule as a whole moving from place to place. (B) Rotational motion is the motion of a turning molecule. (C) Vibrational motion is the back-and-forth movement of a vibrating molecule.

The kinetic energy of a substance is measured as the temperature of that substance.
    • Temperature is actually a measure of the average kinetic energy and has nothing to do with heat until there is a transfer of energy.

The number of oxygen molecules with certain velocities that you might find a sample of air at room at temperature. Notice that a few are barely moving and some have velocities over 1,000 m/s at a given time, but the average velocity is somewhere around 500 m/s.

    • Conceptually a thermometer is used to measure the hotness or coldness of an object.
    • What a thermometer really measures is the average kinetic energy of an object.
    • There is a physical transfer of kinetic energy to the thermometer which them responds due to the increase in its kinetic energy.
      • Mercury.
      • Ethylene glycol.

(A) A bimetallic strip is two different metals, such as iron and brass, bonded together as a single unit, shown here at room temperatures. (B) Since one metal expands more than the other, the strip will bend when it is heated. In this example, the brass expands more than the iron, so the bimetallic strip bends away from the brass.


This thermostat has a coiled bimetallic strip that expands and contracts with changes in the room temperature. The attached vial of mercury is tilted one way or the other, and the mercury completes or breaks an electric circuit that turns the heating or cooling system on or off.

Thermometer Scales.
    • Fahrenheit scale
      • Sets boiling point of water at 212 OF and freezing point of water at 32 OF.
      • 180 divisions between these two.
      • Like most English measures is quite cumbersome.
Celsius scale.
    • Sets boiling point of water at 100 OC and freezing point of water at 0 OC.
    • 100 divisions between these two points.
  • Kelvin or absolute scale.
    • Begins at absolute zero, the temperature at which all kinetic energy is changed into potential energy.
      • ie, all molecular motion ceases.
    • Boiling point of water is 373 K and freezing point of water is 273 K.
    • Divisions are same as for Celsius scale
    • TF = 1.8 TC + 32 OC
    • TC = (TF - 32 OF)
  • 1.8
      • 1.8 accounts for the divisions between freezing point of water and boiling point of water.
        • There are 1.8 divisions in the F scale for every 1 division in the C scale.
    • TK = TC + 273.
    • The temperature of Lake Superior in August averages 34 OF. What is the temperature in OC.
    • Use: TC = (TF - 32 OF)
    • 1.8
    • TC = (34 OF - 32 OF)
    • 1.8
    • TC = (2 OF)
    • 1.8
    • What is the equivalent Celsius temperature of 400.0 K? The equivalent Fahrenheit temperature?
    • Use: TK = TC + 273
    • Rearrange to : TC = TK - 273
    • TC = 400.0 K - 273 = 127.0 OC
    • TF = 1.8 (127.0 OC) + 32 OC =
Internal and External Energy.
    • External energy is total potential and kinetic energy of everyday sized objects.
    • Internal energy is the total kinetic and potential energy of an object molecules.

One theory about how friction results in increased temperatures: Molecules on one moving surface will catch on another surface, stretching the molecular forces that are holding it. They are pulled back to their home position with a snap, resulting in a gain of vibrational kinetic energy.


External energy is the kinetic and potential energy that you can see. Internal energy is the total kinetic and potential energy of molecules. When you push a table across a floor, you do work against friction. Some of the external mechanical energy goes into internal kinetic and potential energy, and the bottom surface of the legs becomes warmer.

Heat as Energy Transfer.
    • Temperature is a measure of the average kinetic energy of an object.
    • Heat is a measure of the internal energy that has been absorbed or transferred from one body to another.
      • Increasing the internal energy is called heating.
      • Decreasing the internal energy is called cooling.
Two ways to increase temperature:
    • From a temperature difference, with energy moving from a region of higher temperature to a region of lower temperature.
    • From an object gaining energy by way of a temperature conversion.

Heat and temperature are different concepts, as shown by a liter of water (1,000 mL) and a 250 mL cup of water, both at the same temperature. You know the liter of water contains more heat since it will require more ice cube to cool it, say, 25OC than will be required for the cup of water. In fact, you will have to remove 48,750 additional calories to cool the liter of water

Measures of Heat.
    • The metric unit of measuring work, energy, or heat is the joule.
    • The metric unit of heat is the calorie.
      • A calorie is the amount of energy needed to increase the temperature of 1 gram of water 1 OC (from 14.5 OC to 15.5 OC.
      • A kilocalorie is the amount of energy needed to increase the temperature of 1 kg of water 1 OC.

The Calorie value of food is determined by measuring the heat released from burning the food. If there is 10.0 kg of water and the temperature increased from 10OC to 20OC the food contained 100 Calories (100,000 calories). The food illustrated here would release much more energy than this.


Joule worked with the English system of measurement used during his time. When a 100 lb object falls 7.78 ft, it can do 778 fl?lb of work. If the work is done against friction, as by stirring 1 lb of water, the heat produced by the wok raises the temperature 1OF.

The English unit of heating is the BTU.
    • A BTU is the amount of energy needed to increase the temperature of 1 lb of water 1 OF.
    • A Quad is 1 quadrillion BTU 1 X 1015 BTU.
    • 778 ftlb = 1 BTU
    • 4.184 ftlb = 1 calorie
    • 4,184 J = 1 kcalorie
Example: a 2,200.0 kg automobile is moving at 90.0 km/hr (25.0 m/s). How many kilocalories are generated when the car brakes to a stop?
    • KE = 1/2mv2
    • KE = 1/2(2,200 kg)(25.0m/s)2
    • KE = (1,100 kg)(625.0 m2/s2)
    • KE = 687,500 m2/s2
    • KE = 687,500 J
    • Kcal = 687,500 J X 1 kcal/4,184J = 164 kcalories
Specific Heat.
    • Three variables that influence energy transfer.
      • The temperature change.
      • The mass of the substance.
      • The nature of the material being heated.
    • The amount of heat (Q) needed to increase the temperature (Ti) of a pot of water from the initial temperature to a final temperature (Tf) is proportional to (Tf-Ti).
      • Q  (Tf-Ti).
      • Q T.
The quantity of heat (Q) absorbed or given off during a certain change in temperature is also proportional to the mass (m) of the substance.
    • Q  m
  • Putting this all together we get:
    • Q  mcT
    • c is the specific heat of the substance.
    • Specific heat is the energy needed to increase the temperature of 1 gram of a substance 1 OC.
When two materials of different temperatures are involved in heat transfer and are perfectly insulated from the surroundings, the heat lost by one is equal to the heat gained by the other.
    • heat lost = heat gained.
    • Qlost = Qgained
    • (mcT)lost = (mcT)gained
Of these three metals, aluminum needs the most heat per gram per degree when warmed, and releases the most heat when cooled.
Example: How much heat must be supplied to a 500.0 g pan to increase its temperature from 20.0 OC to 100.0 OC if the pan is made of a) iron and b) aluminum.
    • Iron from table 5.2 has a specific heat of 0.11 cal/gOC.
    • Q = mcT
    • Q = (500.0g)(0.11 cal/gOC)(80.0OC)
    • Q = 4,400 cal or 4.40 kcalories
    • Aluminum from table 5.2 has a specific heat of 0.22 cal/gOC
    • Q = mcT
    • Q = (500.0g)(0.22 cal/gOC)(80.0OC)
    • 8,800 calories or 8.80 kcalories
Heat Flow.
    • Conduction.
      • Anytime there is a temperature difference; there is a natural tendency for temperature to flow from the area of higher temperature to the area of lower temperature.
      • Conduction is the transfer of energy from molecule to molecule.
      • The rate depends on the temperature difference, the area and thickness of the substance, and the nature of the material.

Thermometers place in holes drilled in a metal rod will show that heat is conducted from a region of higher temperature to a region of lower temperature. The increased molecular activity is passed from molecule to molecule in the process of conduction.

Some materials are good conductors while others are good insulators.
    • Conductors transfer energy very efficiently.
    • Insulators transfer energy very inefficiently,
    • The best conductors are usually metals which have very little air space between molecules.
    • The best insulators have a great deal of air space between molecules.
    • The absolute best insulator is a vacuum as there are no molecules to pass on energy.
Fiberglass insulation is rated in terms of R-value, a ratio of the conductivity of the material to its thickness.
    • Large scale transfer of heat by a large scale displacement of groups of molecules with relatively higher kinetic energy.
    • Molecules with higher kinetic energy are moved from one place to another place.
    • Happens only in liquids and gases where fluid motion can carry molecules with higher kinetic energy over a distance.

(A) Two identical volumes of air are balanced, since they have the same number of molecules and the same mass. (B) Increased temperature causes one volume to expand from the increased kinetic energy of the gas molecules. (C) The same volume of the expanded air now contains fewer gas molecules and is less dense, and it is buoyed up by the cooler, more dense air.


Convection currents move warm air throughout a room as the air over the heater becomes warmed, expands, and is moved upwards by cooler air.

    • Radiation involves the form of energy called radiant energy that moves through space.
    • All objects with a temperature above absolute zero give off radiant energy.
    • The absolute temperature of the object determines the rate, intensity, and kinds of radiant energy emitted.
Phase Change.
    • The motion of a molecule can be increased by:
      • Adding heat through a temperature difference.
      • The absorption of one of the five forms of energy.
      • Temperature increases according to the specific heat of the substance.
When a substance changes from one state to another, the transition is called a phase change.
    • A phase change always absorbs of releases energy, a quantity of heat that is not associated with a temperature change.
    • Latent heat is the hidden energy of a phase change, which is energy that goes in or comes out of internal potential energy.
Three major types of phase change.
    • Solid-liquid.
    • Liquid-gas.
    • Solid-gas
    • The temperature at which a substance changes from a liquid to a solid is called the freezing point.
    • The temperature at which a solid changes to a liquid is the melting point.
    • Both of these occur at the same temperature.
    • The temperature at which a liquid changes from the liquid phase to the gaseous phase is the boiling point.
    • The temperature at which a gas or vapor changes to the liquid phase is the condensation point.
    • Both of these occur at the same temperature.
    • A phase change directly from a solid to a gas or vapor is called sublimation.
Each phase change absorbs or releases a quantity of latent heat, which goes into or is released from molecular potential energy.

This graph shows three warming sequences and two phase changes with a constant input of heat. The ice warms to the melting point, then absorbs heat during the phase change as the temperature remains constant. When all the ice has melted, the now liquid water warms to the boiling point, where the temperature again remains constant as heat is absorbed during the second phase change from liquid to gas. After all the liquid has changed to gas, continued warming increases the temperature of the water vapor.


(A)Work is done against gravity to lift an object, giving the object more gravitational potential energy. (B) Work is done against intermolecular forces in separating a molecule from a solid, giving the molecule more potential energy.


Compare this graph to the one in Figure 5.20. This graph shows the relationships between the quantity of heat absorbed during warming and phase changes as water is warmed from ice at -20OC to water vapor at some temperature above 100OC. Note that the specific heat for ice, liquid water, and water vapor (steam) have different values.

Latent heat of fusion.
    • The latent heat of fusion is the heat involved in a solid-liquid phase change in melting or freezing.
    • A melting solid absorbs energy and a freezing liquid releases this same amount of energy, warming the surroundings.
    • The total heat involved in a solid-liquid phase change depends on the mass of the substance involved.
      • Q = mLf
      • Where Lf is the latent heat of fusion for the substance involved
Latent heat of vaporization.
    • The amount of heat involved during a phase change from a liquid to a gas or vapor is called the latent heat of vaporization.
    • The latent heat of vaporization is the heat involved in a liquid-gas phase change where there is evaporation or condensation.
    • The escaping molecules absorb energy from the surroundings, and a condensing gas releases this exact same amount of energy.
The total heating depends on the amount of water vapor condenses so that:
    • Q = mLV
    • Where LV is the latent heat of vaporization for the substance involved.
    • How much energy does a refrigerator remove from 100.0 g of water at 20.0 OC to make ice at -10.0 OC
    • Three steps.
    • Q1 = mcT to cool from 20.0 OC to 0.0 OC
    • Q1 = (100.0g)(1.00cal/gOC)(0.0OC-20.0OC)
    • = 2,000 cal = 2.00 X 10 3 cal.
Q2 = mLf to remove latent heat of fusion.
  • (100.0g)(80.0cal/g)
  • 8,000 cal = 8.00 X 103 cal
  • Q3 = mcT to go from 0.0 OC to -10 OC
  • (100.0g)(0.500cal/g)(10.0OC-0.0OC)
  • 500 cal = 5.00 X 102 cal
  • Qtotal = Q1 + Q2 + Q3
  • = 2.00 X 103 cal + 8.00 X 103 cal + 5.00 X 102 cal
  • = 1.05 X 104 cal
Evaporation and Condensation.
    • Evaporation occurs when enough energy is inputed into a system to cause liquid molecules to overcome attractive forces near the surface, escape, and become a gas or vapor.
    • In evaporation, more molecules are leaving the liquid state than are returning.
    • In condensation, more molecules are returning to the liquid state than are leaving.
    • When the condensation rate is equal to the evaporation rate, the air above the liquid is saturated (holds all the vapor that it is capable of holding).

Temperature is associated with the average energy of the molecules of a substance. These numbered circles represent arbitrary levels of molecular kinetic energy that, in turn, represent temperature. The two molecules with the higher kinetic values [25 in (A)] escape, which lowers the average value from 11.5 to 8.1 (B). Thus evaporation of water molecules with more kinetic energy contributes to the cooling effect of evaporation in addition to the absorption of latent heat.

Four ways to increase the rate of evaporation.
    • An increase in temperature of the liquid will increase the average kinetic energy of the molecules and thus increase the number of high energy molecules capable of escaping from the liquid state.
    • Increase the surface area of the liquid in contact with the air.
Removal of water vapor from near the surface will prevent the return of molecules to the liquid phase.
  • Reducing atmospheric pressure will reduce one of the forces holding molecules in a liquid.
Relative Humidity.
    • The ratio of how much water vapor is in air to how much water vapor it could hold at a certain temperature is the relative humidity
    • Usually expressed as a percent.

The inside of this closed bottle is isolated from the environment so the space above the liquid becomes saturated. While it is saturated, the evaporation rate equals the condensation rate. When the bottle is cooled, condensation exceeds evaporation and droplets of liquid form on the inside surfaces.

The curve shows the maximum amount of water vapor in g/m3 that can be in the air at various temperatures.
    • The laws of thermodynamics describe what happens to energy as it is transformed into work and to other forms.
    • Thermodynamics is concerned with internal energy, which is the total internal kinetic and potential energy of a system.
The system is the component we want to describe.
  • The state of the system are the variable under which it exists, temperature, pressure, volume, heat, etc…
  • Everything outside of the system is the surroundings.

A very simple heat engine. The air in (B) has been heated, increasing the molecular motion and thus the pressure. Some of the heat is transferred to the increased gravitational potential energy of the weight as it is converted to mechanical energy.

The First Law of Thermodynamics.
    • The energy supplied to a system is equal to the change in internal energy
The Second Law of Thermodynamics.
    • Heat flows from objects with a higher temperature to objects with a cooler temperature.
The Second Law and Natural Processes.
    • Energy can be viewed from two considerations of scale:
      • The observable external energy of an object.
      • The internal energy of the molecules, or particles that make up an object.
Two kinds of motion that the particles of an object can have.
    • A coherent motion where they move together.
    • An incoherent, chaotic motion of individual particles.
  • Work on an object is associated with coherent motion, while heating an object is associated with its internal incoherent motion.

The heat supplied (QH) to a heat engine goes into the mechanical work (W) and the remainder is expelled in the exhaust (QL). The work accomplished us therefore the difference in the heat input and output (QH) - (QL), so the work accomplished represents the heat used, W = J(QH - QL)


A heat pump uses work (W) to move heat from a low temperature region (QL) to a high temperature region (QH). The heat moved (QL) requires work (W), so J QL = W.

    • Energy is always degrading toward a more disorderly state.
    • The total entropy of the universe is continually increasing.
    • The natural process is for the sate of order to degrade into a state of disorder with a corresponding increase in entropy.
Eventually all of the useable energy in the universe will diminish to unusable forms.
    • The universe will at some time reach a limit of disorder called the heat death of the universe.
    • The heat death of the universe is the theoretical limit of disorder, with all molecules spread far, far apart, vibrating slowly with a uniform low temperature.