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Chapter 4. Heat and Temperature. Matter. All matter is made up of tiny, basic units called atoms. Atoms cannot be created, destroyed, or divided during chemical or physical changes.

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Chapter 4

Chapter 4

Heat and Temperature


  • All matter is made up of tiny, basic units called atoms.

  • Atoms cannot be created, destroyed, or divided during chemical or physical changes.

  • Each element in the periodic table is made up of a different type of atom. e.g. hydrogen, oxygen, carbon, iron, gold.

  • Compounds are made up of more than one element. e. g. water, sugar, alcohol.


  • Temperature is a measure of heat content of atoms and molecules.

  • The higher the temperature the more heat the atoms and/or molecules of the substance contain.

  • The higher the heat and temperature the higher the average kinetic energy of the atoms and molecules of the substance (element or compound).

  • The higher the average kinetic energy of the atoms and molecules of the substance the faster they move, since KE=1/2 mv2.

  • The movement can be translation (straight line, occurs only in gases), vibration, or rotation (occurs in gases, liquids, and solids).

A Solid-Vibration only

B Liquid-Vibration and Rotation only

C Gas-Vibration, Rotation and Translation

Temperature scales
Temperature Scales

  • There are 3 temperature scales that are used:

    Fahrenheit (oF)

    Celsius or Centigrade(oC)

    Kelvin (Absolute Temperature)(K)

Conversions between temp scales
Conversions between Temp Scales

  • 1o Celsius is approximately 1.8oF, or 9/5 oF.

  • K = °C + 273

  • °C = (°F – 32)


  • °F = (°C x 1.8) + 32

  • Boiling Point of Water is 100oC, 212oF, 373K

  • Freezing Point of Water is 0oC, 32oF, 273K

Absolute temperature
Absolute Temperature

  • Kelvin is the absolute temperature.

  • At 0K all movement of molecules and atoms stops.

  • This is the SI scale of temperature.

  • 0K = -273oC or -459.4oF

  • 0K has never been reached, but scientists have come to 700 billionths of a Kelvin above it.

Temperature conversions
Temperature Conversions

  • Convert 34oC to oF:

    oF = (oC x 1.8) + 32

    oF = (34oC x 1.8) + 32 = 93oF

  • Convert – 22.7oF to oC:

    oC = (oF – 32)


    oC = (-22.7oF – 32) = -30.4oC


  • Convert -30.4oC to K:

    K = oC + 273

    K = -30.4oC + 273 = 242.6 K

External and internal energy
External and Internal Energy

  • Potential and Kinetic Energy

    • For External Energy, PE and KE is same as chapter 3

    • For Internal Energy, PE and KE applies to molecules of substance.


  • Measure of internal energy absorbed or transferred from one body to another.

    • Process of increasing internal energy is heating, decreasing is cooling.

    • Heat is always transferred from a hotter object to a colder object.

    • Eventually both objects will be at the same temperature.

Two heating methods
Two Heating Methods require more ice

  • From temperature difference, with energy moving from the region of higher temperature to the region of lower temperature.

  • From an object gaining energy by way of an energy form conversion, from mechanical or electrical energy to heating; etc.

Measuring heat
Measuring Heat require more ice

  • The unit for energy is the Joule (J) and the Kilojoule (kJ) in the metric system.

  • The unit that is sometimes used for heat, however is the calorie or kilocalorie (1000 calories).

  • A calorie is the amount of heat which is needed to increase the temperature of 1 gram of water by 1oC. The kilocalorie is the amount of heat needed to increase the temperature of 1 kg of water by 1oC.

  • In the English system the heat unit is the BTU (British thermal unit). This one is the amount of heat needed to increase the temperature on 1 pound of water by 1oF.

Conversion calorie and joule
Conversion calorie and joule require more ice

  • 1 cal = 4.184 J

  • I kcal = 1000 cal

  • 1 kcal = 4.184 kJ

Specific heat
Specific Heat require more ice

  • Q = m x c x ΔT

    Q = heat

    m = mass

    c = specific heat

    ΔT = Change in Temperature

    c, specific heat, is a physical property associated with

    every substance. It depends on the internal structure of the substance (the atoms and molecules)

    The value for c indicates how much heat is required to raise the temperature of 1 g of a substance by 1oC.

    The units for specific heat are cal/goC or J/goC

    The specific heat for water (c water) is 1 cal/goC

Specific heat1
Specific heat require more ice

  • Specific heat is inherent to each individual substance.

Specific heat2
Specific Heat require more ice

  • How much heat is absorbed by 57.3 g of a substance that has a specific heat of 2.7 cal/goC if the temperature increases by 5.7 oC ?

    Q = ?Q = c x m x ΔT

    C= 2.7 cal/goCQ = 57.3 g x 2.7 cal/goC x 5.7 oC

    ΔT = 5.7 oCQ = 881.8 cal

    m = 57.3 g

Specific heat3
Specific heat require more ice

  • Example:

    • 5000 cal of heat is added to 100 g of water at 20°C, what is the final temperature?

      Q=5000 cal Q = m x c x ΔT

      m = 100 g ΔT = Q

      m x c

      T1 = 20°C

      ΔT = 5000 cal

      100 g x 1 cal/goC

      T2 = ?

      C = 1 cal/g°C

      ΔT = 50 oC

      ΔT = T2 – T1

      T2 = ΔT+ T1

      T2=50 oC + 20°C

      T2= 70°C

In coastal areas temperature vary less than inland. require more ice

Because water has a higher specific heat than land,

it ends up absorbing much of the atmospheric heat,

which helps keep temperatures cooler. Water releases heat gradually

when the atmosphere is cold, which keeps temperatures warmer.

Phase changes
Phase Changes require more ice

  • The phases of matter are solid, liquid, and gas.

  • When heated substances increase in temperature no matter what phase they are in, but when a phase change is occurring the temperature remains constant.

  • Heat added during a phase change is called latent heat and it refers to the energy that comes into or comes out of internal potential energy.

Phase changes1
Phase Changes require more ice

  • 3 major phase changes and they can go in either direction:

  • Solid to liquid: forward is melting (or fusion), reverse is freezing

  • Liquid to gas: forward is vaporization (includes evaporation and boiling), reverse is condensation

  • Solid to gas: forward is sublimation, reverse is deposition.

Phase changes2
Phase Changes require more ice

  • Solid to Liquid occurs at the melting point.

  • Liquid to Solid occurs at the freezing point.

  • The melting point and the freezing point are the same temperature.

  • Liquid to Gas occurs at the boiling point.

  • Gas to Liquid occurs at the condensation point.

  • The boiling point and the condensation point occur at the same temperature.

Phase changes3
Phase Changes require more ice

  • Solid to Gas occurs at the sublimation point.

  • Gas to Solid occurs at the deposition point.

  • The sublimation point and the deposition point occur at the same temperature.

  • Dry ice sublimes to gaseous carbon dioxide and mothballs, which are made of naphthalene, sublime to the gaseous state at standard atmospheric pressure.

Latent heat
Latent Heat require more ice

Energy to turn water from a liquid to

a gas or vice versa!

Energy to turn water from a solid to a liquid or vice versa!

Lv is the latent heat of vaporization, per gram of the substance.

Lf is the latent heat of fusion, per gram of the substance.

Evaporation substance.

  • Liquids can evaporate at temperatures lower than their boiling point when left out in the open. e.g. a water puddle.

  • There are particles of liquid that have more kinetic energy than others and these have enough energy to escape into the gas phase.

  • Once these particles escape another batch will take their place with enough kinetic energy to escape and eventually they will all evaporate.

After evaporation occurs the remaining substance is at a cooler

temperature. This is until the substance is again warmed by the

atmosphere and the process repeats.

Perspiration in humans is a cooling process.

Thermodynamics cooler

  • The study of heat and its relationship to mechanical energy.

  • Concerned with the internal energy, U, the total internal potential and kinetic energies of molecules making up a substance.

Laws of thermodynamics
Laws of Thermodynamics cooler

  • The first law of thermodynamics:

    Energy can neither be created nor destroyed. It can only be changed from one form to another.

Laws of thermodynamics1
Laws of Thermodynamics cooler

  • As a result: the energy supplied to a thermodynamic system in the form of heat, (Qin) minus the work done by the system, (W) is equal to to the change in internal energy of the system (Qout).

  • W = Qin- Qout

Heat: Lost energy

Second law of thermodynamics
Second Law of Thermodynamics cooler

  • Heat flows from objects with a higher temperature to objects with a lower temperature.

    • The total entropy of any isolated thermodynamic system tends to increase over time, approaching a maximum value.

Entropy cooler

  • The total entropy of the universe continually increases.

  • Entropy is a thermodynamic measure of disorder.

  • Order means patterns and coherent arrangements. Disorder means dispersion, no patterns, and a randomized or spread-out arrangement.

  • Entropy is a measure of chaos

Third law of thermodynamics
Third Law of Thermodynamics cooler

  • As a system approaches absolute zero of temperature (0K) all processes virtually cease and the entropy of the system approaches a minimum value; also stated as: "the entropy of a perfectly crystalline body at absolute zero temperature is zero."

Homework chapter 4
Homework Chapter 4 cooler

  • Applying Concepts, p. 111-112, # 1, 3, 4, 5, 6, 7, 8, 9, 10, 15, 16, 17, 18, 21

  • Parallel Exercises, p. 112-113 # 1, 2, 7, 9

  • New Book:

  • p. 118-120 # 3, 5, 6, 11, 12, 13, 14, 15, 16, 18, 23, 24, 26, 27, 28, 29, 30, 31, 32, 40, 41

  • Parallel Exercises, p. 121 #1, 2, 3, 7, 9.

Review chapter 4

Matter, atoms, molecules, elements, compounds. cooler

Temperature-What is it (A measure of the amount of heat a body contains)

Relationship among temperature, heat, kinetic energy (The higher the temperature the higher the heat content and the kinetic energy content)

Types of movement in solids, liquids and gases. (Translation, vibration and rotation in gases; only vibration and rotation in liquids and solids).

Temperature scales.(oF, oC, K)

Temperature conversions

Freezing and boiling points of water

Internal (atoms and molecules) and External Energy (normal size particles)

Specific heat-Amount of energy in calories or joules needed to increase the temperature of 1 g of a substance by 1oC.

Calorie & Kilocalorie

1 cal = 4.184 J

Specific Heat

Q = m x c x ΔT

Specific heat of water (1 cal/goC)

Phase changes-freezing and melting (s-l); condensation and vaporization (l-g); deposition and sublimation (s-g).

s-l, l-g, and s-g are endothermic. l-s, g-l, and g-s are exothermic. The temperature remains constant during all phase changes.

fp and mp; cp and bp; dp and sp.

Fusion is the same as melting.

Vaporization includes evaporation (below the boiling point) and boiling.

Latent heat of vaporization and fusion (Amount of energy needed to convert a liquid to a gas and a solid to a liquid).

Diagrams for phase changes

Evaporation-Molecules at the surface with enough kinetic energy escape to the vapor state.

Laws of Thermodynamics:

1o Law-Energy cannot be created or destroyed. Energy added is converted to work + internal energy.

2o Law-Heat is transferred from a warmer object to a colder object. The entropy of the universe is always increasing.

3o Law-All movement stops at 0K (Absolute 0).

Entropy-State of disorder.

Review Chapter 4