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During a
“physical change”
a substance changes
some physical
property…
H2O
…but it is still the
same material with
the same chemical
composition.
gas
H2O
solid
liquid
Chemical Property:
The tendency of a
substance to change into
another substance.
Steel rusting:
caused by iron (Fe)
reacting with oxygen (O2)
to produce rust (Fe2O3)
4 Fe + 3 O2
2 Fe2O3
Chemical Change:
Any change involving a
rearrangement of atoms.
Chemical Reaction:
The process of a
chemical change...
During a
“chemical reaction”
new materials are
formed by a change
in the way atoms are
bonded together.
Physical and Chemical Properties
Examples of Physical Properties
Boiling point
Color
Slipperiness
Electrical conductivity
Melting point
Taste
Odor
Dissolves in water
Shininess (luster)
Softness
Ductility
Viscosity (resistance to flow)
Volatility
Hardness
Malleability
Density (mass / volume ratio)
Examples of Chemical Properties
Burns in air
Reacts with certain acids
Decomposes when heated
Explodes
Reacts with certain metals
Reacts with certain nonmetals
Tarnishes
Reacts with water
Is toxic
Ralph A. Burns, Fundamentals of Chemistry 1999, page 23
Chemical properties can ONLY be observed during a chemical reaction!
The formation of a
compound
The formation of a
mixture
Physical & Chemical Changes
CO2
crushing
heating
Pyrex
PHYSICAL
CHANGE
Limestone,
CaCO3
CHEMICAL
CHANGE
CaO
Crushed limestone,
CaCO3
Lime and
carbon dioxide,
CaO + CO2
Sunlight
energy
O2
Pyrex
Pyrex
H2O2
H2O
Light hastens the decomposition of hydrogen peroxide, H2O2. H
O
The dark bottle in which hydrogen peroxide is usually stored
keeps out the light, thus protecting the H2O2 from decomposition.
H
O
Properties of Ionic Compounds
•
•
•
•
•
•
•
•
Crystalline solids
Hard and brittle
High melting points
High boiling points
High heats of vaporization
High heats of fusion
Good conductors of electricity when molten
Poor conductors of heat and electricity when
solid
• Many are soluble in water
Chemical Bonds
• between two identical nonmetal atoms are non-polar covalent.
• between two different nonmetal atoms are polar covalent.
• between nonmetals and reactive metals are primarily ionic.
Covalent bonding
Electrons are shared
equally
Cl
Cl
Polar covalent bonding
Electrons are shared
unequally
H Cl
Increasing ionic character
Ralph A. Burns, Fundamentals of Chemistry 1999, page 229
Ionic bonding
Electrons are
transferred
Na1+
Cl1-
Chemical Bonds
• between two identical nonmetal atoms are nonpolar covalent.
• between two different nonmetal atoms are polar covalent.
• between nonmetals and reactive metals are primarily ionic.
Nonpolar covalent
Polar covalent
Electrons are shared
equally
Electrons are shared
unequally
Cl
Cl
H Cl
Increasing ionic character
Ralph A. Burns, Fundamentals of Chemistry 1999, page 229
Ionic bonding
Electrons are
transferred
Na1+
Cl1-
Covalent vs. Ionic
Alike
Different
Share
electrons
Different
Transfer
electrons
Chemical
Bonds
(polar vs. nonpolar)
(ions formed)
+/-
Topic
Between
Two
Nonmetals
Covalent
Topic
Electrons
are
involved
Ionic
Between
Metal and
Nonmetal
Weak
Bonds
Strong
Bonds
(low melting point)
(high melting point)
Celsius & Kelvin Temperature Scales
Kelvin
Celsius
Boiling point
of water
Freezing point
of water
Absolute
zero
100oC
100
Celsius
degrees
373 K
100
Kelvins
0oC
273 K
-273oC
0K
Temperature is Average Kinetic Energy
“HOT”
Fast
“COLD”
Slow
Kinetic Energy (KE) = ½ m v 2
*Vector = gives direction and magnitude
Temperature Scales
Boiling point
of water
Fahrenheit
Celcius
Kelvin
212 oF
100 oC
373 K
180 oF
Freezing point
of water
32 oF
100 oC
0 oC
100 K
273 K
Notice that 1 kelvin degree = 1 degree Celcius
Kelvin Scale
blue
white
yellow
5250 K
Sunlight
5000 K
PIAA Plasma Blue
4300 K
PIAA HID Bulb
4150 K
PIAA Xtreme White
3800 K
PIAA Super White
3200 K
Halogen Bulb
2600 K
Incandescent Bulb
Temperature Scales
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 136
Compare Celsius to Fahrenheit
oF
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 139
– 32 = 1.8 oC
Converting
70 degrees
Celsius to
Kelvin
units.
oC
+ 273 = K
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 137
Heat versus Temperature
lower temperature
Fractions of particles
higher temperature
TOTAL
= Heat
Kinetic ENERGY
Kinetic energy
Molecular Velocities
molecules sorted by speed
Fractions of particles
many different molecular speeds
the Maxwell speed distribution
speed
http://antoine.frostburg.edu/chem/senese/101/gases/slides/sld016.htm
Temperature vs. Heat
Different
Alike
Measured
with a
Thermometer
Have
Kinetic
Energy
Topic
Average
Kinetic
Energy
oCelcius
(or Kelvin)
Temperature
Different
Measured
with a
Calorimeter
Topic
A Property
of
Matter
Heat
Total
Kinetic
Energy
Joules
(calories)
Conservation of Matter
Reactants
yield
Products
Density
•
Density is an
INTENSIVE property
of matter.
- does NOT depend
Brick
Styrofoam
on quantity of
matter.
- color, melting point, boiling point, odor, density
•
Contrast with
EXTENSIVE
- depends on
quantity of matter.
- mass, volume, heat content (calories)
Properties of Matter
Pyrex
Pyrex
Extensive
Properties
volume:
mass:
100 mL
99.9347 g
15 mL
14.9902 g
Intensive
Properties
density:
temperature:
0.999 g/mL
20oC
0.999 g/mL
20oC
http://antoine.frostburg.edu/chem/senese/101/matter/slides/sld001.htm
Styrofoam
D =
Styrofoam
M
V
Brick
D
=
Brick
M
V
Volume and Density
Relationship Between Volume and Density for Identical Masses of Common Substances
Substance
Cube of substance
(face shown actual size)
Mass
(g)
Volume
(cm3)
19
Density
(g/cm3)
Lithium
10
0.53
Water
10
10
1.0
Aluminum
10
3.7
2.7
Lead
10
0.58
11.4
Density
M
D =
V
M
M = DxV
ass
D
ensity
V
olume
M
V =
D
Density of Some
Common Substance
Density of Some Common Substances
Substance
Density
(g / cm3)
Air
Lithium
Ice
Water
Aluminum
Iron
Lead
Gold
*at 0oC and 1 atm pressure
0.0013*
0.53
0.917
1.00
2.70
7.86
11.4
19.3
Density of Carbon Dioxide
Carbon
Dioxide
CO2
Density Air = 1.29 g/L
Density CO2 = 1.96 g/L
Carbon
Dioxide
CO2
Symptoms of CO Poisoning
Concentration of CO
in air (ppm)*
Hemoglobin
molecules as HbCO
100 for 1 hour or less
10% or less
500 for 1 hour or less
20%
500 for an extended period
of time
30 - 50%
headache, confusion, nausea,
dizziness, muscular weakness,
fainting
1000 for 1 hour or less
50 - 80%
coma, convulsions, respiratory
failure, death
*ppm is parts per million
Davis, Metcalfe, Williams, Castka, Modern Chemistry, 1999, page 760
Visible effects
no visible symptoms
mild to throbbing headache,
some dizziness, impaired
perception
Carbon Monoxide Poisoning
‘The Silent Killer’
Hemoglobin (Hb) binds with carbon monoxide (CO) in the capillaries of the lungs.
Poisoning: Hb + CO  HbCO
If caught in time, giving pure oxygen (O2) revives victim of CO poisoning.
Treatment causes carboxyhemoglobin (HbCO) to be converted slowly to
oxyhemoglobin (HbO2).
Treatment: O2 + HbCO  CO + HbO2
Carbon monoxide, CO, has almost 200 times the affinity to bind
with hemoglobin, Hb, in the blood as does oxygen, O2.
Davis, Metcalfe, Williams, Castka, Modern Chemistry, 1999, page 760
Exchange of Blood Gases
Tank of Water
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 143
Person Submerged in Water
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 143
Archimedes Principle
Thread
Vfinal = 98.5 cm3
- Vinitial = 44.5 cm3
Vfishing sinker = 54.0 cm3
98.5 cm3
44.5 cm3
Fishing sinker
Water
Before immersion
After immersion
Galilean Thermometer
• Density = Mass / Volume
• Mass is constant
• Volume changes with temperature
Temp = 68 oC
– Increase temperature  larger volume
In the Galilean thermometer, the small glass bulbs are partly
filled with a different (colored) liquid. Each is filled with a
slightly different amount, ranging from lightest at the
uppermost bulb to heaviest at the lowermost bulb. The clear
liquid in which the bulbs are submerged is not water, but
some inert hydrocarbon (probably chosen because its
density varies with temperature more than that of water
does).
Galilean Thermometer
In the Galilean thermometer, the small
glass bulbs are partly filled with a different
(colored) liquid. Each is filled with a slightly
different amount, ranging from lightest at
the uppermost bulb to heaviest at the
lowermost bulb. The clear liquid in which
the bulbs are submerged is not water, but
some inert hydrocarbon (probably chosen
because its density varies with temperature
more than that of water does).
The correct temperature is the lowest
floating bulb. As temperature increases,
density of the clear medium decreases
(and bulbs sink).
RECALL: Density equals mass / volume.
76oF
80o
68 oF
80o
76o
76o
72o
72o
68o
68o
64o
64o
Solid, Liquid, Gas
(a) Particles in solid
(b) Particles in liquid
(c) Particles in gas
Solid
H2O(s) Ice
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 31
Ice
Photograph of snowflakes
Photograph of ice model
H2O(s) Ice
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
In a liquid
• molecules are in
constant motion
Liquid
• there are appreciable
intermolecular forces
• molecules are close
together
• Liquids are almost
incompressible
• Liquids do not fill the
container
H2O(l) Water
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 31
Gas
H2O(g) Steam
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 31
Liquids
The two key properties we need to describe are
EVAPORATION and its opposite CONDENSATION
add energy and break intermolecular bonds
EVAPORATION
CONDENSATION
release energy and form intermolecular bonds
States of Matter
Gas, Liquid, and Solid
Gas
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 441
Liquid
Solid
States of Matter
Solid
heat
Liquid
heat
Gas
Holds Shape
Shape of Container
Shape of Container
Fixed Volume
Free Surface
Volume of Container
Fixed Volume
Some Properties of Solids, Liquids, and Gases
Property
Solid
Liquid
Gas
Shape
Has definite shape
Takes the shape of
the container
Takes the shape
of its container
Volume
Has a definite volume
Has a definite volume Fills the volume of
the container
Arrangement of
Particles
Fixed, very close
Random, close
Random, far apart
Interactions between
particles
Very strong
Strong
Essentially none
Evaporation
• To evaporate, molecules must have
sufficient energy to break IM forces.
• Molecules at the surface break away
and become gas.
• Only those with enough KE escape.
• Breaking IM forces requires energy. The
process of evaporation is endothermic.
• Evaporation is a cooling process.
• It requires heat.
Condensation
Change from gas to liquid
Achieves a dynamic equilibrium with
vaporization in a closed system.
What is a closed system?
A closed system means
matter can’t go in or out.
(put a cork in it)
What the heck is a
“dynamic equilibrium?”
Dynamic Equilibrium
When first sealed, the molecules
gradually escape the surface of the
liquid.
As the molecules build up above the
liquid - some condense back to a
liquid.
The rate at which the molecules
evaporate and condense are equal.
Dynamic Equilibrium
As time goes by the rate of vaporization
remains constant but the rate of
condensation increases because there
are more molecules to condense.
Equilibrium is reached when:
Rate of Vaporization = Rate of Condensation
Molecules are constantly changing phase “dynamic”
The total amount of liquid and vapor remains constant
“equilibrium”
Vaporization
• Vaporization is an endothermic process - it
requires heat.
• Energy is required to overcome intermolecular
forces
• Responsible for cool earth
• Why we sweat
Energy Changes Accompanying Phase Changes
Gas
Energy of system
Vaporization
Condensation
Sublimation
Liquid
Melting
Freezing
Solid
Brown, LeMay, Bursten, Chemistry 2000, page 405
Deposition
MATTER
yes
MIXTURE
yes
Is the composition
uniform?
Homogeneous
Mixture
(solution)
PURE SUBSTANCE
no
Heterogeneous
Mixture
Colloids
no
Can it be physically
separated?
yes
Can it be chemically
decomposed?
Compound
Suspensions
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
no
Element
Both elements and compounds have a definite makeup and definite properties.
Elements
only one kind
of atom; atoms
are bonded it
the element
is diatomic or
polyatomic
substance
with
definite
makeup
and
properties
Packard, Jacobs, Marshall, Chemistry Pearson AGS Globe, page (Figure 2.4.1)
Compounds
two or
more kinds
of atoms
that are
bonded
Mixtures
two or
more
kinds of
and
two or more
substances
that are
physically
mixed
Matter Flowchart
Examples:
– graphite
element
– pepper
hetero. mixture
– sugar (sucrose)
compound
– paint
hetero. mixture
– soda
solution
homo. mixture
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Pure Substances
Element
– composed of identical atoms
– EX: copper wire, aluminum foil
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Pure Substances
Compound
– composed of 2 or more elements
in a fixed ratio
– properties differ from those of
individual elements
– EX: table salt (NaCl)
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Pure Substances
Law of Definite Composition
– A given compound always contains the same,
fixed ratio of elements.
Law of Multiple Proportions
– Elements can combine in different ratios to
form different compounds.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Pure Substances
For example…
Carbon, C
Carbon, C
Oxygen, O
Oxygen, O
Oxygen, O
Carbon monoxide, CO
Carbon dioxide, CO2
Two different compounds,
each has a definite composition.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Mixtures
Variable combination of two or more
pure substances.
Heterogeneous
Homogeneous
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Mixtures
Solution
– homogeneous
– very small particles
– no Tyndall effect
Tyndall Effect
– particles don’t settle
– EX: rubbing alcohol
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Mixtures
Colloid
– heterogeneous
– medium-sized particles
– Tyndall effect
– particles don’t settle
– EX: milk
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Mixtures
Suspension
– heterogeneous
– large particles
– Tyndall effect
– particles settle
– EX: fresh-squeezed
lemonade
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Mixtures
Examples:
– mayonnaise
colloid
– muddy water
suspension
– fog
colloid
– saltwater
solution
– Italian salad
dressing
suspension
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Classification of Matter
Materials
Homogeneous
Heterogeneous
Substance
Element
Compound
Homogeneous
mixture
Heterogeneous
mixture
Solution
Order / Disorder
Smoot, Smith, Price, Chemistry A Modern Course, 1990, page 43
Mixture
Classification of Matter
MATTER
(gas. Liquid,
solid, plasma)
Separated by
PURE
SUBSTANCES
MIXTURES
physical means into
Separated by
COMPOUNDS
ELEMENTS
chemical
means into
Kotz & Treichel, Chemistry & Chemical Reactivity, 3rd Edition , 1996, page 31
HOMOGENEOUS
MIXTURES
HETEROGENEOUS
MIXTURE
Classification of Matter
uniform
properties?
fixed
composition?
no
heterogeneous
mixture
no
solution
no
element
yes
compound
chemically
decomposable?
http://antoine.frostburg.edu/chem/senese/101/matter/slides/sld003.htm
Elements, Compounds, and Mixtures
hydrogen
atoms
oxygen atoms
(a)
an element
(hydrogen)
(b)
a compound
(water)
hydrogen
atoms
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 68
(c)
a mixture
(hydrogen
and oxygen)
(d)
a mixture
(hydrogen
and oxygen)
Elements, Compounds, and Mixtures
hydrogen
atoms
oxygen atoms
(a)
an element
(hydrogen)
(b)
a compound
(water)
hydrogen
atoms
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 68
(c)
a mixture
(hydrogen
and oxygen)
(d)
a mixture
(hydrogen
and oxygen)
Mixture vs. Compound
Different
Alike
Variable
Composition
Involve
substances
Topic
No bonds
between
components
Can be
separated by
physical means
Mixture
Different
Fixed
Composition
Topic
Contain
two or more
elements
Can be
separated
into
elements
Compound
Bonds
between
components
Can ONLY be
separated by
chemical means
Compounds vs. Mixtures
• Compounds have properties that are
uniquely different from the elements from
which they are made.
– A formula can always be written for a compound
– e.g. NaCl  Na + Cl2
• Mixtures retain their individual properties.
– e.g. Salt water is salty and wet
Diatomic Elements, 1 and 7
H2
N2 O2 F2
Cl2
Br2
F2
The Haber Process
Matter
Physically
separable
Substance
Definite composition
(homogeneous)
Element
(Examples: iron, sulfur,
carbon, hydrogen,
oxygen, silver)
Chemically
separable
Mixture of
Substances
Variable composition
Compound
(Examples: water.
iron (II) sulfide, methane,
Aluminum silicate)
Homogeneous mixture
Heterogeneous mixture
Uniform throughout,
also called a solution
(Examples: air, tap water,
gold alloy)
Nonuniform
distinct phases
(Examples: soup,
concrete, granite)
The Organization of Matter
MATTER
HOMOGENEOUS
MIXTURES
HETEROGENEOUS
MIXTURE
Physical methods
PURE
SUBSTANCES
ELEMENTS
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 41
Chemical
methods
COMPOUNDS
Top Ten Elements
in the Universe
Element
1. Hydrogen
2. Helium
3. Oxygen
4. Carbon
5. Neon
6. Iron
7. Nitrogen
8. Silicon
9. Magnesium
10. Sulfur
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 26
Percent
(by atoms)
73.9
24.0
1.1
0.46
0.13
0.11
0.097
0.065
0.058
0.044
A typical spiral galaxy
(Milky Way is a spiral galaxy)
Properties of Matter
•
•
•
•
•
•
•
Electrical Conductivity
Heat Conductivity
Density
Melting Point
Boiling Point
Malleability
Ductility
Paper Chromatography
Paper Chromatography
of Water-Soluble Dyes
orange
red
Filter paper
(stationary phase)
yellow
Suggested
red dye
is not
homogeneous
Orange mixture of
red
and
yellow
Initial spots of dyes
Direction of Water
(mobile phase)
movement
Specific Heats
of Some Substances
Specific Heat
Substance (cal/ g oC)
Water
Alcohol
Wood
Aluminum
Sand
Iron
Copper
Silver
Gold
1.00
0.58
0.42
0.22
0.19
0.11
0.093
0.057
0.031
(J/g oC)
4.18
2.4
1.8
0.90
0.79
0.46
0.39
0.24
0.13
Potential energy
Energy in
kinetic energy
Energy out
kinetic energy
The energy something possesses due to its motion, depending on mass and velocity.
Energy
A
C
B
Kinetic Energy – energy of motion
KE = ½ m v 2
mass
velocity (speed)
Potential Energy – stored energy
Batteries (chemical potential energy)
Spring in a watch (mechanical potential energy)
Water trapped above a dam (gravitational potential energy)
School Bus or Bullet?
Which has more kinetic energy;
a slow moving school bus or a fast moving bullet?
Recall: KE = ½ m v 2
BUS
KE = ½ m v 2
KE(bus) = ½ (10,000 lbs) (0.5 mph)2
BULLET
KE = ½ m v 2
KE(bullet) = ½ (0.002 lbs) (240 mph)2
Either may have more KE, it depends on the mass of the bus and the velocity
of the bullet.
Which is a more important factor: mass or velocity? Why?
(Velocity)2
Kinetic Energy and Reaction
Rate
lower temperature
Fractions of particles
higher temperature
minimum energy
for reaction
Kinetic energy
Kinetic Energy and Reaction
Rate
lower temperature
Fractions of particles
higher temperature
minimum energy
for reaction
Kinetic energy
Hot vs. Cold Tea
Many molecules have an
intermediate kinetic energy
Low temperature
(iced tea)
High temperature
(hot tea)
Percent of molecules
Few molecules have a
very high kinetic energy
Kinetic energy
Exothermic Reaction
Reactants  Products + Energy
10 energy
=
8 energy
+ 2 energy
Energy of reactants
Energy
Energy of products
Reactants
-DH
Products
Reaction Progress
Endothermic Reaction
Energy + Reactants  Products
Energy
Activation
Energy
Reactants
Products
+DH Endothermic
Reaction progress
Effect of Catalyst on Reaction Rate
WhatCatalyst
is a catalyst?
does it do
duringfor
a chemical
reaction?
lowers What
the activation
energy
the reaction.
No catalyst
Energy
activation energy
for catalyzed reaction
reactants
products
Reaction Progress
An Energy Diagram
activated
complex
Ea
activation
energy
energy
reactants
products
course of reaction
Animation by Raymond Chang
All rights reserved