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.
• 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
• 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
• There are 3 temperature scales that are used:
Fahrenheit (oF)
Celsius or Centigrade(oC)
Kelvin (Absolute Temperature)(K)
Conversions between Temp Scales
• 1o Celsius is approximately 1.8oF, or 9/5 oF.
• K = °C + 273
• °C = (°F – 32)
1.8
• °F = (°C x 1.8) + 32
• Boiling Point of Water is 100oC, 212oF, 373K
• Freezing Point of Water is 0oC, 32oF, 273K
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
• 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)
1.8
oC = (-22.7oF – 32) = -30.4oC
1.8
• Convert -30.4oC to K:
K = oC + 273
K = -30.4oC + 273 = 242.6 K
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.
Heat
• 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.
The liter of water contains more internal energy. It would require more ice
cubes to cool it.
Two Heating Methods
• 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
• 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
• 1 cal = 4.184 J
• I kcal = 1000 cal
• 1 kcal = 4.184 kJ
Specific Heat
• 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 heat
• Specific heat is inherent to each individual
substance.
Specific Heat
• 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=?
C= 2.7 cal/goC
ΔT = 5.7 oC
m = 57.3 g
Q = c x m x ΔT
Q = 57.3 g x 2.7 cal/goC x 5.7 oC
Q = 881.8 cal
•
Specific heat
Example:
– 5000 cal of heat is added to 100 g of water at 20°C, what is the final temperature?
Q=5000 cal
m = 100 g
T1 = 20°C
Q = m x c x ΔT
ΔT = Q
mxc
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.
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
• 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 Changes
• 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 Changes
• 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 Changes
• 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
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
• 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
• 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
• The first law of thermodynamics:
Energy can neither be created nor
destroyed. It can only be changed from
one form to another.
Laws of Thermodynamics
• 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
• 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
• 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 spreadout arrangement.
• Entropy is a measure of chaos
Third Law of Thermodynamics
•
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
• 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
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Matter, atoms, molecules, elements, compounds.
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
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Specific heat of water (1 cal/goC)
Phase changes-freezing and melting (s-l);
condensation and vaporization (l-g);
deposition and sublimation (s-g).
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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.
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fp and mp; cp and bp; dp and sp.
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Fusion is the same as melting.
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Vaporization includes evaporation (below
the boiling point) and boiling.
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Latent heat of vaporization and fusion
(Amount of energy needed to convert a
liquid to a gas and a solid to a liquid).
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Diagrams for phase changes
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Evaporation-Molecules at the surface with
enough kinetic energy escape to the vapor
state.
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Laws of Thermodynamics:
o
1 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).
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Entropy-State of disorder.