States of Matter Solid Has definite shape & volume
Particles are tightly packed Can expand when heated
Slide 5
States of Matter Liquid Has constant volume but takes the shape
of its container Fluid Less closely packed particles than solid
particles Can expand when heated
Slide 6
States of Matter Gas Expands to fill its container & takes
the shape of its container Fluid Much less closely packed than
solid particles Expands when heated
Slide 7
Forces of Attraction There are two kinds of attractive forces
at the molecular level The forces inside a molecule holding the
individual atoms together The forces between molecules holding
different molecules together in a sample
Slide 8
Intramolecular Forces Intra- prefix = within The forces inside
a molecule holding the individual atoms together Ex.) Covalent
bonds in H 2 O
Slide 9
Demo What happened to the paper clip when placed in the beaker
of water vs. the beaker of acetone? Explain your observations.
Slide 10
Intermolecular Forces Inter- prefix = between Short range
forces between molecules in a sample There are 3 main types of
intermolecular forces Hydrogen bonding Dipole-dipole forces London
Dispersion forces
Slide 11
Dipole-Dipole Forces
Slide 12
Dipole a molecule or part of a molecule that contains both
positively and negatively charged regions + (partial positive) or -
(partial negative) Dipole-Dipole Forces forces of attraction
between POLAR molecules Dipoles must get close together in correct
orientation (positive end must be near negative end)
Slide 13
Dipole-Dipole Forces H-FH-FH-FH-F Dipole-dipole forces will
raise melting and boiling points. A dipole can temporarily attract
electrons from another molecule causing an induced dipole.
Slide 14
London Dispersion Forces
Slide 15
Intermolecular attractions resulting from the uneven
distribution of electrons and the creation of temporary dipoles
Present in all substances (polar molecules, non- polar molecules,
and noble gases) The weakest intermolecular force Does the number
of electrons play a role regarding strength of the
attraction??
Slide 16
London Dispersion Forces Electrons are constantly moving around
the nucleus therefore electron density can fluctuate This effect
becomes stronger with increasing number of electrons Example: F 2
Br 2 I 2
Slide 17
Hydrogen Bonding
Slide 18
Hydrogen Bonding attractive forces in which a hydrogen
covalently bonded to a very electronegative atom (F, O, N) is also
weakly bonded to an unshared electron pair on another
electronegative atom (another F, O, or N atom) Hydrogen bonds are
strong intermolecular forces
Slide 19
Hydrogen Bonding Hydrogen bonding raises melting and boiling
points because more energy is required to break the forces between
molecules.
Slide 20
Hydrogen Bonding in DNA Hydrogen bonding plays a key role in
maintaining the double helix structure of DNA
Slide 21
Think Back to the Demo Question Why did the paperclip float on
water but not acetone? Water Acetone
Slide 22
Before You Go: Identify the types of intermolecular forces
present in compounds of: Hydrogen Fluoride Pentane (C 5 H 12 )
Hydrochloric Acid Ethanol (Ethyl Alcohol)
Slide 23
Relative Strength of Intermolecular Forces (Strongest)
(Weakest)
Slide 24
Liquids & Solids Even though they are more restricted than
gas molecules, the molecules of solids and liquids are constantly
in motion as well! (This idea comes from the Kinetic Molecular
Theory well come back to this idea)
Slide 25
Liquids Viscosity a measure of the resistance of a liquid to
flow The particles in a liquid are close enough together that their
attractive forces slow their movement as they flow past one another
The stronger the attractive forces (intermolecular forces), the
more viscous the liquid is. As temperature increases, viscosity
decreases.
Slide 26
Surface tension an inward force that tends to minimize the
surface area of a liquid A measure of the inward pull by particles
in the interior The stronger the intermolecular forces, the higher
the surface tension Liquids In water, this is due mainly to
hydrogen bonding!
Slide 27
Liquids Surfactant any substance that interferes with the
hydrogen bonding between water molecules & reduces surface
tension
Slide 28
Surfactants used to clean up oil spills as well Exxon Valdez
oil spill in 1989 spilled over 700,000 barrels of oil into the
water near Alaska
Slide 29
Solids Crystalline solid a solid in which the atoms, ions, or
molecules are arranged in an orderly, geometric, three-dimensional
structure Unit cell the smallest arrangement of connected points
that can be repeated to form the lattice A.k.a. The representative
part of the whole crystal
Network Covalent Solids Atoms that can form multiple covalent
bonds (such as C, Si, and other Group 14 elements) are able to form
network covalent solids. All atoms in the entire structure are
bonded together with covalent chemical bonds.
Slide 33
Metallic Solids Metallic solids positive metal ions surrounded
by a sea of mobile electrons Mobile electrons make metals malleable
and ductile because electrons can shift while still keeping the
metal ions bonded in their new places Metallic solids are good
conductors of heat and electricity Metallic Bonds
Slide 34
Amorphous Solids Amorphous solid a solid in which the particles
are not arranged in a regular, repeating pattern Amorphous =
without shape Often form when a molten material cools too quickly
to allow enough time for crystals to form Common examples: glass,
rubber, many plastics
Slide 35
Kinetic Molecular Theory Describes the behavior of gases (and
solids/liquid to some extent) in terms of particles in motion
Assumptions: 1. Gas particles have negligible volume compared to
the volume of their container 2. Particles move in constant,
random, straight line motion 3. Particles collide with themselves
and walls without losing energy (elastic collisions) 4. There are
no intermolecular forces between gas molecules
Slide 36
Kinetic Molecular Theory If gas particles are always in this
constant, random motion, what keeps them going? ENERGY!! (Kinetic
Energy to be exact) Temperature is a measure of the average kinetic
energy in a substance. Higher temp. = more kinetic energy =
particles move faster!
Slide 37
Temperature 3 Main Temperature Scales Fahrenheit Celsius (C) C
= K - 273 Kelvin (K) K = C + 273
Slide 38
Temperature Conversions Convert the following temperatures: 28
C = ________ K 200 K = ________ C -15 C = _________ K
Slide 39
Gases Remember: Gases expand to fill their container & take
the shape of their container In other words, gases will spread out
until they cant spread out anymore Gases will also move according
to diffusion and effusion.
Slide 40
Racing Gases Demo: If concentrated HCl is at one end of the
tube and concentrated NH 3 is at the other end, which gas do you
think will move farthest and fastest down the tube? Racing Gases
Demo HCl (g) NH 3 (g)
Slide 41
RACING GASES DEMO The gases will diffuse down the tube
Diffusion tendency of molecules to move from areas of higher
concentration towards areas of lower concentration Example:
spraying perfume and smelling it across the room
Slide 42
DIFFUSION Originally Over Time
Slide 43
RACING GASES DEMO The gases diffused at different rates If the
white ring forms closer to the HCl end of the tube, which gas moved
farthest and fastest? Why did this gas move faster? BECAUSE ITS
LIGHTER!! (Has a lower molar mass)
Slide 44
Grahams Law of Effusion The racing gases demo is related to
Grahams Law Grahams Law of Effusion gases of lower molar masses
effuse (and diffuse) faster than gases with higher molar masses
Effusion when a gas escapes (diffuses) through a tiny hole in its
container Example: Helium balloons shrinking compared to normal
balloons
Slide 45
Grahams Law Grahams Law be applied to both the effusion and the
diffusion of a gas Gases with lower molar masses (lighter gases)
diffuse faster than gases with higher molar masses (heavier gases)
The lighter the gas, the faster it moves!!!
Slide 46
Grahams Law Which gas would diffuse and effuse faster Methane
(CH 4 ) or carbon dioxide (CO 2 )? Chlorine (Cl 2 ) or oxygen (O 2
)? Hydrogen sulfide (H 2 S) or carbon monoxide (CO)?
Slide 47
Grahams Law Formula The rates are simply the speed or velocity
at which the gas is traveling. So, this formula will compare the
speed of one gas to the speed of the other gas.
Slide 48
Pressure of a Gas The force of a gas exerted on the surface of
a container As gases bounce around in a container, each collision
with the a container wall exerts a force More collisions = higher
pressure Less collisions = lower pressure Empty space with no
particles and no pressure is called a vacuum
Slide 49
Pressure of a Gas Atmospheric pressure collision of air
molecules with objects As elevation increases, air density and
therefore pressure decrease Barometers measure atmospheric
pressure
Slide 50
Pressure of a Gas Vapor pressure pressure due to force of gas
particles above a liquid colliding with walls of container Higher
temp. = higher vapor pressure
Slide 51
Pressure of a Gas Gravity When baking, there are different
instructions for baking at high altitudeswhy? Boiling point of
water decreases since lower pressure Liquid leaves faster so food
must bake longer
Slide 52
Pressure of a Gas SI unit of pressure: pascal (Pa) Other common
pressure units: Millimeters of mercury (mm Hg) Atmospheres (atm) 1
atm = 760 mmHg = 101.3 kPa = 760 torr
Slide 53
Practice Converting Units 1 atm = 760 mmHg = 101.3 kPa A tire
pressure gauge records a pressure of 450 kPa. What is the pressure
in atmospheres? In mm Hg?
Slide 54
DALTONS LAW
Slide 55
Partial pressure the contribution of each gas in a mixture to
the total pressure Daltons Law of Partial Pressures for a mixture
of gases, the total pressure is the sum of the partial pressure of
each gas in the mixture P total = P 1 + P 2 + P 3 + (at constant
volume and temperature) Daltons Law
Slide 56
Practice with Daltons Law Determine the total pressure of a gas
mixture that contains oxygen, nitrogen, and helium. The partial
pressures are: P O 2 = 20.0 kPa, P N 2 = 46.7 kPa, and P He = 26.7
kPa. P total = P 1 + P 2 + P 3 +
Slide 57
Practice with Daltons Law Air contains O 2, N 2, CO 2, and
trace amount of other gases. What is the partial pressure of oxygen
(P O 2 ) if the total pressure of the system is 101.3 kPa and the
partial pressures of N 2, CO 2, and the other gases are 79.10 kPa,
0.040 kPa, and 0.94 kPa, respectively? P total = P 1 + P 2 + P 3
+
Slide 58
Phase Changes What are phase changes? A change in a substances
state of matter What are some examples of phase changes?
Slide 59
Phase Changes that Require Energy If you have to put energy
into a reaction to make it happen, it is called an endothermic
reaction. Endothermic Phase Changes: Melting (solid liquid)a.k.a
fusion Vaporization (liquid gas) Sublimation (solid gas)
Slide 60
Phase Changes that Release Energy If energy is released or
given off by a reaction, it is called an exothermic reaction.
Exothermic Phase Changes: Condensation (gas liquid) Deposition (gas
solid) Freezing (liquid solid)