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HEAT AND TEMPERATURE
Energy: the ability to do work, cause
change or transfer heat
KE = kinetic energy = energy of motion =
½ mv2
PE = potential energy = energy stored in
chemical bonds; cannot be
measured directly
CHAPTER 10 –
MODERN CHEMISTRY TEXT STATES OF MATTER
The Kinetic Molecular Theory of Matter
Laws of Thermodynamics and Heat Transfer
Special Properties of Water
Phase Changes
HEAT ≠ TEMPERATURE
Temperature
Intensive physical property
Heat
= Thermal Energy
≈ Enthalpy
Extensive physical property
Does not depend on size of
sample
Depends on size of sample
Defined: A measure of the
average KE of the particles
in a sample of matter
Defined: a form of energy; the
total E content of system=
stored PE in bonds + KE due to
particle motion
Defined: hotness/coldness
property that controls the
direction of heat flow
Defined: transfer of thermal
energy due to a temperature
difference
Unit = Kelvin (K)
Unit = joules (J)
HEAT ≠ TEMPERATURE
Temperature Units
Heat Units
1 calorie= 4.184 joules
K = °C + 273
°C = K – 273
Absolute Zero = 0 Kelvin
(1 cal = 4.184 J)
1000 joules = 1 kilojoule
(1000 J = 1 kJ )
1 calorie = heat energy needed to
increase temp of 1 gram of water by
1 degree Celsius
1000 calories = 1 Calorie = 1
kilocalorie
(1000 cal = 1 Cal = 1 kcal)
HEAT TRANSFER
First Law of Thermodynamics:
 Energy can be converted from one form to another but
the total energy of a system is conserved.
 Second Law of Thermodynamics:
 Heat flows naturally from hot objects to cooler ones
until a thermal equilibrium is reached.
 Combined laws → rule of heat transfer:
 Heat lost by system = Heat gained by surroundings.

A
simple calorimeter:
The “system” pulls heat
from or transfers heat
to a known amt. of H2O
in a closed rxn. vessel
HEAT (THERMAL ENERGY) TRANSFER
– THE BIG IDEAS




The temperature of a material is a measure of the average kinetic
energy of the molecules that make up that material. Absolute zero is
defined as the temperature at which the molecules have zero kinetic
energy, which is why it is impossible for anything to be colder.
Solids are rigid because their particles do not have enough kinetic
energy to go anywhere—they just vibrate in place. The particles in a
liquid have enough energy to move around one another—which is why
liquids flow—but not enough to escape each other. (Note: In
evaporation, some particles will gain enough KE to escape randomly
from the surface and become a gas. If the liquid is volatile, this
happens faster.) In a gas, the particles have so much kinetic energy
that they disperse and the gas expands to fill its container.
Heat is a measure of how much thermal energy is transmitted from one
body to another. While both work and heat can be measured in joules,
they are not measures of energy but rather of energy transfer!
Source: spark notes heat and temperature
Chapter 10 – Section 1: The Kinetic-Molecular Theory of Matter
THE KINETIC-MOLECULAR THEORY
The kinetic-molecular theory of matter
states:




Particles of matter (atoms and
molecules) are always in motion.
We measure this energy of motion
(kinetic energy) as temperature.
If temperature increases, the
particles will gain more energy
and move even faster.
Molecular motion is greatest in
gases, less in liquids, and least
in solids.
THE KINETIC MOLECULAR THEORY OF MATTER
Explains the properties of solids, liquids, & gases
in terms of …..

the energy (motion) of the particles
in competition with

the attractive forces that act between the particles
(ionic, covalent & metallic bonds and intermolecular
forces).
Chapter 10 – Section 1: The Kinetic-Molecular Theory of Matter
Gases
• An Ideal Gas is a hypothetical gas that
perfectly fits all the assumptions of the kineticmolecular theory.
• Many gases behave nearly
ideally if pressure is not very
high and temperature is not
very low.
• Fluidity – Gas particles glide easily past one
another. Because liquids and gases flow, they
are both referred to as fluids.
Chapter 10 – Section 1: The Kinetic-Molecular Theory of Matter
Gases (continued)
• Low Density – Gas particles are very far apart. The
density of a gas is about 1/1000 the density of the
same substance in the liquid or solid state.
• Expansion – A gas will expand to fill its container.
• Compressibility – The volume of a gas can be greatly
decreased by pushing the particles closer together.
Chapter 10 – Section 1: The Kinetic-Molecular Theory of Matter
Gases (continued)
• Diffusion– spontaneous mixing of two gases due to
random motion and empty space between particles.
THE KINETIC MOLECULAR THEORY OF GASES
FIVE ASSUMPTIONS ABOUT IDEAL GASES:
1.
2.
3.
4.
5.
Gases are composed of a large number of particles in
a state of constant, random motion; these particles
collide with other particles and the walls of the
container.
These particles are much smaller than the distance
between particles. Most of the volume of a gas is
therefore empty space.
There is no force of attraction between gas particles.
Collisions between gas particles or collisions with
the container walls are perfectly elastic - None of
the KE of a gas particle is lost.
The average KE of a gas depends on the
temperature of the gas and nothing else.
REAL GASES VS IDEAL GASES




Real Gases ≈ Ideal Gases IF the
temperature is not too low AND the
pressure is not too high
High pressure and low temperature…
will cause real gases to have attractions
between particles!!!
Noble gases are most likely to behave like
ideal gases under most (ordinary)
conditions of temperature and pressure
(because they are not polar)
Polar gas molecules are least likely to
behave like ideal gases (because they have
dipoles which causes attractions).
Chapter 10 – Section 2: Liquids
Liquids
• Surface Tension – Strong
cohesive forces at a liquid’s
surface act to decrease the
surface area to the smallest
possible size. The higher the
force of attraction between
the particles of a liquid, the
higher the surface tension.
Chapter 10 – Section 2: Liquids
Liquids (continued)
• Vaporization – A liquid or solid
changing to a gas.


Evaporation – particles escape
from the surface of a liquid and
become a gas. This occurs
because liquid particles have
different kinetic energies.
Boiling – bubbles of vapor appear throughout a liquid.
Will not occur below a certain temperature (the boiling
point.)
• A volatile liquid is one that evaporates readily.
• A vapor is the gaseous state of a volatile liquid;
signifies that both phases are present.
SPECIAL PROPERTIES OF WATER
 Water
is a bent, very polar molecule with
Hydrogen bonds between them:
 Water
bonds:
has high cohesion because of H
SPECIAL PROPERTIES OF WATER
 Water is a great solvent for other polar
substances and for ionic compounds:
SPECIAL PROPERTIES OF WATER

But not for oils and fats which are nonpolar
SPECIAL PROPERTIES OF WATER
Ice Floats! Most matter is more dense as a
solid than a liquid but water reaches max
density at 4°C and expands when it freezes!
Transient
H bonds
Stable H Bonds
WATER RESISTS TEMPERATURE CHANGES
GOOD FOR YOU & THE CLIMATE
1. high specific heat capacity: it takes a lot of energy to
heat water
2. high heat of vaporization: you have to add a lot of energy
to change state to a gas… steam stores a lot of energy as
heat which is released as it condenses AND absorbed when
it evaporates = cooling effect on your body
3. high heat of fusion: you have to remove a lot of energy to
get it to freeze…
WHY?
H bonds require a lot of
energy to break,
minimizing temp changes
SPECIAL PROPERTIES OF WATER
Ice Floats! Most matter is more dense as a
solid than a liquid but water reaches max
density at 4°C and expands when it freezes!


https://youtu.be/45yabrnryXk
http://astrobiology.nasa.gov/articles/life-withoutwater/
Chapter 10 – Section 3: Solids
Solids
• There are two main types of solids:


Crystalline Solids – Made up
of crystals. Particles are
arranged in an orderly,
geometric, repeating pattern, sharp MPt.
Amorphous Solid – Particles
are arranged randomly, no sharp MPt.
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
Ionic Bonding and the Crystal Lattice
• In an ionic crystal, ions minimize their potential
energy by combining in an orderly arrangement
known as a crystal lattice.
• A formula unit is the smallest repeating unit of an
ionic compound.
Sodium Chloride crystal lattice (many Na and Cl atoms)
Formula Unit = NaCl
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
Comparing Ionic and Covalent Compounds
• Covalent compounds have relatively weak forces of
attraction between molecules, but ionic compounds
have a strong attraction between ions. This causes
some differences in their properties:
Ionic
crystals
very high melting points
hard, but brittle
Ex: NaCl, CaF2, KNO3
Covalent
molecules
low melting points
usually gas or liquid
Ex: H2O, CO2, O2
Chapter 6 – Section 4: Metallic Bonding
The Metallic Bond
• In metals, overlapping orbitals allow the outer
electrons of the atoms to roam freely
throughout the entire metal.
• These mobile electrons form
a sea of electrons around the
metal atoms, which are packed
together in a crystal lattice.
• A metallic bond results from the attraction
between metal atoms and the surrounding
sea of electrons.
Chapter 10 – Section 3: Solids
Solids (continued)
• Melting Point – The temperature at which a
solid becomes a liquid. At this temperature, the
kinetic energies of the particles within the solid
overcome the attractive forces holding them
together.
SUMMARY




SOLIDS: attractions > motion
LIQUIDS: attractions ≈ motion
GASES: attractions < motion
can change phase solid→ liguid → gas
only if motion (KE) overcomes attractions
PHASE CHANGES
PHASE CHANGES
HEATING CURVE
HEATING CURVES
Sloped portions = Δ temp = Δ KE only
Level portions = Δ of state = Δ PE only
LAWS OF THERMODYNAMICS
1.
2.
3.
4.
5.
Heat flows “downhill” from object with higher temp. to
object with lower temp. until equilibrium is reached.
No heat flows b/w two bodies at same temp.
Absolute zero = zero Kelvin, lowest possible temp. in
nature, all molecular motion ceases.
When energy is converted to work, the process is never
100% efficient. The graveyard of all lost energy is heat.
Energy changes form but is conserved in all chemical
and physical changes.
THE THREE LAWS OF
THERMODYNAMICS

The first law, also known as Law of Conservation
of Energy, states that energy cannot be created or
destroyed in an isolated system.The second law
of thermodynamics states that the entropy of any
isolated system always increases.The third law of
thermodynamics states that the entropy of a system
approaches a constant value as
the temperature approaches absolute zero.
Source: Boundless. “The Three Laws of
Thermodynamics.” Boundless Chemistry. Boundless,
03 Feb. 2016. Retrieved 10 Mar. 2016
from https://www.boundless.com/chemistry/textbooks/
boundless-chemistry-textbook/thermodynamics17/the-laws-of-thermodynamics-123/the-three-laws-ofthermodynamics-496-3601/
STATES OF MATTER & BOND TYPE
 Metallic
atoms → metallic bonding → solids
 Metallic
cations with nonmettalic anions→
ionic bonding → solids
 Two
or more nonmetallic atoms → covalent
bonding →individual molecules →
 weak attractions b/w molecules,
intermolecular forces, determine phase
 sometimes solid, usually liquid OR gas
INTERMOLECULAR FORCES
Defined: Weak forces of
attraction b/w molecules.
Compared to bonds, (intra
molecular forces), they are much
weaker.
Without IMF, all molecular
substances would be gases
including water!
If IMF ↑ particles have greater
attractions b/w them
BPts, MPts, surface tension
and viscosity also ↑
evaporation rates &
volatility will ↓
INTERMOLECULAR FORCES
WEAKEST TO STRONGEST
 Dispersion
Forces:
 Force of attraction b/w two nonpolar
molecules with temporary dipoles
 Dipole-Dipole forces:

the force of attraction b/w two polar
molecules
 Hydrogen

Bonding:
Force of attraction b/w hydrogen (low EN)
and fluorine, oxygen, and nitrogen – three
nonmetals with high EN values.
TYPES OF SOLIDS
Crystalline Solids
Metallic Solids
 Covalent (molecular) Solids
 Ionic Solids
 Covalent Network Crystals

Amorphous Solids:
“Without shape”
 Atoms are NOT arranged in a regular pattern
 EX: Glass and plastic
 No sharp MPt

CRYSTALLINE SOLIDS
Single crystals or groups of crystals fused together.
 The 3-D arrangement has a coordinate system called a
lattice.
 The smallest unit of the lattice is a unit cell.
 Type of crystal based on:



Particles in the crystal
Types of bonding b/w the particles
PROPERTIES OF SOLIDS







Definite shape and volume
Highly ordered, low KE
Not compressible, rigidly retains shape &
volume
High density, particles are close together
Extremely low diffusion and expansion
Tightly packed particles can only vibrate
from fixed positions
ATTRACTIONS (ionic bonds, metallic bonds
or strong intermolecular forces) > MOTION
PROPERTIES OF LIQUIDS








Definite volume and indefinite shape, takes
shape of container
Medium disorder, medium KE
Not very compressible
High density, particles are close together
Slow diffusion and low expansion
Particles can flow past each other, so “fluid”
Exhibit properties of surface tension &
cappilary action.
ATTRACTIONS (intermolecular forces) ≈
DISRUPTIONS CAUSED BY MOTION
PROPERTIES OF LIQUIDS
Surface tension: a force that tends to pull adjacent parts
of a liquid’s surface together, decreasing surface area to a
minimum.
EX: a raindrop
PROPERTIES OF LIQUIDS

Capillary Action: the attraction of the surface
of a liquid to the surface of a solid, tends to pull
the liquid molecules upward along the surface
and against the pull of gravity.
PROPERTIES OF LIQUIDS

Adhesion: water attracted to other materials


EX: meniscus
Cohesion: water attracted to itself

EX: a drop of water on a glass slide
PROPERTIES OF GASES







Indefinite volume & shape, takes shape and volume of
container, fills container, also a “fluid”
Very disordered, very high KE
Very compressible
Very low density, particles are far apart
Rapid diffusion, effusion and expansion
Rapid, random particle movement causes collisions with
other particles and with container… causes gas pressure
ATTRACTIONS < DISRUPTIONS CAUSED BY MOTION
PROPERTIES OF GASES
Expansion – will fill container
 Fluidity – no attractive forces
 Low density & compressibility – particles far apart
 Diffusion: spontaneous mixing of the particles of two
substances due to rapid, random motion and empty
space b/w particles
 Effusion: gas particles pass through tiny openings;
gases with low masses effuse faster
 A vapor is the gaseous state of a volatile liquid signifies that both phases are present.

SOME CHARACTERISTICS OF GASES, LIQUIDS AND SOLIDS AND
THE MICROSCOPIC EXPLANATION FOR THE BEHAVIOR
gas
liquid
solid
assumes the shape and volume of
its container
no definite shape or volume
disordered, high KE
assumes the shape of the part of
the container which it occupies
definite volume but indefinite
shape
medium disorder, medium KE
retains a fixed volume and shape
rigid - particles locked into place in
a regular pattern
highly ordered,
low KE
compressible
not easily compressible
not easily compressible
lots of free space between particles little free space between particles
little free space between particles
rapid diffusion & expansion
limited diffusion and low expansion very limited diffusion and very low
expansion
flows easily
particles can move past one
another &
particles move rapidly and
randomly in all directions
flows easily
particles remain close together but
can move/slide past one another
does not flow easily
rigid - particles can only vibrate in
a fixed position