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Matter and Change
Chapter 1
Matter
anything that has mass and
takes up space
pure substances
– compounds
and elements
mixtures – two or
more pure
substances
mixed in the
same
container…not
bonded
Solid, Liquid, or Gas?
SHAPE
• Solid
• definite shape
VOLUME
• definite volume
• Liquid • indefinite shape
• definite volume
• Gas
• indefinite volume
• indefinite shape
Physical Properties
• physical property – a characteristic that
can be observed or measured without
changing the identity of the substance
melting point
mass
state of matter
color
Physical Changes
• physical change – a change in a
substance that does not involve a change
in the identity of the substance
breaking
cutting
dissolving
boiling
tearing
All phase changes
are physical changes.
Chemical Properties
• chemical property – a characteristic that
can be observed or measured with a
change in the identity of the substance
flammability
reacts with an acid
reacts with oxygen
Chemical Changes
• chemical change – a change in a
substance that does involve a change in
the identity of the substance
color change
gas released
(often with an odor)
energy change
(light, heat, …)
precipitate
formed
…four good indications of a chemical change.
Precipitate
aqueous –
dissolved
in water
• precipitate – a solid formed from two
aqueous solutions during a chemical
reaction
Homogeneous Mixtures
the same throughout…each
sample contains the same
ratio of ingredients
Heterogeneous Mixtures
different throughout…each
sample contains a different
ratio of ingredients
density
Separation of Mixtures
PHYSICAL changes only
magnetism
by hand
evaporation
filtration
chromatography
Separation of elemental Fe
magnetism
This doesn’t work with
iron that is bound into
a compound, only with
elemental Fe.
Distillation Apparatus
Homogeneous Mixture,
Heterogeneous Mixture or
Pure Substance?
Observations vs. Interpretations
• observations – the facts…what you see or
measure
Example: The solution turned cloudy white.
The test tube felt warm to the touch.
• interpretations – your opinion of what you
see
Example: From the previous observations, a
chemical change has occurred.
Qualitative vs. Quantitative
Observations
• Qualitative observations – ones that do not
involve numbers
• Quantitative observations – ones that DO
use numbers (quantity)
Qualitative vs. Quantitative
Which one is it?
1. A white precipitate was formed.
1. qualitative
2. The reaction produced 40.0 g of water.
2. quantitative
3. The test tube felt warm.
3. qualitative
4. The temperature rose 25 degrees.
4. quantitative
Extensive vs. Intensive Properties
• extensive – DOES depend on the amount
of matter present
Example: mass, volume, amount of energy
within a substance
• intensive – does NOT depend on the
amount of matter present
Examples: color, melting point, density, luster
If mass and volume are extensive properties, why
is density an intensive property?
Chapter 2
Measurements and Calculations
Pages 28-65
Prefix
Symbol
Exponential
tera
T
1012
giga
G
109
mega
M
106
kilo
k
103
hecto
h
102
deka
D or da
101
-----
-----
100
deci
d
10-1
centi
c
10-2
milli
m
10-3
micro
m
10-6
nano
n
10-9
pico
p
10-12
Quantity
Quantity
Symbol
Unit Name
Unit
abbreviation
Length
l
meter
m
Mass
m
kilogram
kg
Time
t
second
s
Temperature
T
kelvin
K
Amount of a Substance
n
mole
mol
Electric Current
I
ampere
A
Luminous Intensity
Iv
candela
cd
Size of Units
1L=1dm3
1cm3 = 1mL
Density
Mass
Density
M
D=
V
Density is an INTENSIVE property…
it does NOT depend on the amount of
matter you have.
Volume
1. What is the density of a block of marble that has the
dimensions 5.00cm length, 4.00cm width, and 15.5cm
height and has a mass of 853 g?
M
D=
V
853 g
D=
5.00 cm x 4.00 cm x 15.5 cm
g
D = 2.75
3
cm
2. Diamond has a density of 3.26 g/cm3. What is the mass
of a diamond that has a volume of 0.351 cm3?
M
D=
V
g
M
3.26
=
3
3
cm
0.351 cm

g 
3
0.351
cm
 3.26

3
cm 

M = 1.14 g


 M
3. What is the volume of a sample of liquid mercury that
has a mass of 76.2 g, given that the density of mercury is
13.6 g/mL?
M
D=
V
g
76.2 g
13.6
=
mL
V
76.2 g
V=
g
13.6
mL
V = 5.60 mL
Density water
1.00 g
g
=
= 1.00
1.00 mL
mL
Significant Figures
• digit or figure – 0,1,2,3,4,5,6,7,8,9
• significant digit or figure – a digit that helps
you to understand the details of the entire
number given to you
Significant Figures
5005 m
0.0045 L
4.500 g
100 cm3
100. cm3
0.04020 g
Scientific Notation
• Move the decimal until one non-zero digit
appears to the left of the decimal.
• Be sure the power of the ten reflects the
direction of that move.
• Keep the same number of significant
figures in the scientific notation as your
original.
50050 m = 5.005 x 104 m
0.00450 L = 4.50 x 10-3 L
Accuracy
(getting all
measurements
right)
Precision
(getting all
measurements
the same)
Percent Error
Percentage error =
Valueexperimental - Valueaccepted
Valueaccepted
x 100%
What is the percent error for a mass measurement of
17.7 g, given that the correct value is 21.2 g?
17.7 g - 21.2 g
%E =
x 100%
21.2 g
%E = - 17%
In order to prepare for the chemistry exam
on Wednesday, you should study for at
least 1.0 hour the night before the exam.
How many milliseconds is this?
3.6 x 106 ms
The summer Olympics showed
us that Usain Bolt could run
10.36 meters per second. How
fast is this in miles per hour?
(3.281 ft = 1.000 meters)
23.18 miles/hr
The volume of the Powerade® that I drank
yesterday was 946 mL. What is that
volume in dm3?
.946 dm3 or 9.46 x 10-1 dm3
Chapter 3: Atoms
The Building Blocks of Matter
An atom is the smallest particle of an element that
retains the chemical properties of that element.
The Early Atom
• As early as 400 B.C., Democritus
called nature’s basic particle the
“atomon” based on the Greek word
meaning “indivisible”.
• Aristotle succeeded Democritus
and did not believe in atoms.
Instead, he thought that all matter
was continuous. It was his theory
that was accepted for the next 2000
years. (Read page 43 of your
textbook.)
Basic Laws of Matter
• Law of Conservation of Mass- mass is
neither created nor destroyed during
ordinary chemical reactions or physical
changes.
CH4 + 2O2 → 2H2O + CO2
16g + 64g → 36g + 44g
Antoine Lavoisier
stated this about 1785
Basic Laws of Matter
• Law of Definite Proportions – no matter how much
salt you have, it is always 39.34% Na and 60.66% Cl by
mass.
Example: Sodium chloride always contains
39.34% Na and 60.66% Cl by mass.
2NaCl
100g
116.88g
→ 2Na + Cl2
→ 39.34g + 60.66g
→
? + ?
Joseph Louis Proust
stated this in 1794.
Basic Laws of Matter
• Law of Multiple Proportions- Two or more
elements can combine to form different
compounds in whole-number ratios.
Example
John Dalton
proposed this
in 1803.
Dalton’s Atomic Theory
• In 1808, Dalton proposed a theory to
summarize and explain the laws of
conservation of mass, definite proportions,
& multiple proportions.
I was a school
teacher at the
age of 12!
Dalton’s Atomic Theory
John Dalton - 1808
1. All matter is composed of extremely small particles
called atoms.
2. Atoms of a given element are identical in size,
mass, and other properties.**
3. Atoms cannot be subdivided, created, or
destroyed.**
4. Atoms of different elements combine in simple
whole-number ratios to form chemical compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged.
**Today, we know these parts to have flaws.
Flaws of Dalton’s Theory…
2. Atoms of a given element are
identical in size, mass, and other
properties. Isotopes – atoms with the same
number of protons but a different
number of neutrons
3. Atoms cannot be subdivided, created,
or destroyed.
Subatomic particles – electrons,
protons, neutrons, and more
The Atom
• Atom - the smallest particle of an element
that retains the chemical properties of that
element.
CARBON
Subatomic Particles
• Protons- positively charged particles found in
the nucleus of an atom.
• Neutrons- neutral particles found in the nucleus
of an atom.
• Electrons- negatively charged particles found in
the electron cloud.
Joseph John Thomson
• In 1897 the English physicist Joseph John
Thomson was able to measure the ratio of
charge of the cathode ray particles to their mass.
• He found that the ratio was always the same
regardless of the metal used to make the
cathode or the nature of the gas inside the
cathode ray tube.
• Thomson concluded that cathode rays were
composed of identical, negatively charged
particles called electrons.
Cathode Ray Tube Experiment
Accomplishments
• Proved that the atom was divisible and that
all atoms contain electrons.
• This contradicted Dalton’s Atomic Theory.
• This allowed a new model of the atom.
Discovery of the Nucleus
• In 1911, Ernest
Rutherford performed a
Gold Foil Experiment.
• He and his colleagues
bombarded a thin piece
of gold foil with fast
moving, positively
charged alpha particles.
Alpha Particles
• Alpha (a) particles are Helium-4 nuclei.
• This means they are two protons and two
neutrons (with no electrons).
• Thus, they are positive.
4
2
He
+2
Gold Foil Experiment
Gold Foil
Experiment
Gold Foil Experiment
• The volume of the
nucleus was very small
compared to the volume
of the atom.
• Therefore, most of the
atom was composed of
empty space. Niels Bohr
later found that this empty
space was where the
electrons were located.
Bohr’s Model of The Atom
Atomic Number
atomic number (Z) - the number of protons
in the nucleus of each atom of a given
element. – (Henry Moseley)
The number of p+ identifies the element.
Atomic Number increases from left to right
on the periodic table.
Electrons
The number of electrons in a neutral atom is
equal to the number of protons in that atom.
e- = p+
•Electrons can be lost or gained.
• When electrons are lost or gained, ions are
formed.
Metals vs. Nonmetals
• Metals form cations.
Na  Na+ + 1e• Nonmetals form anions.
Cl + 1 eCl-

Mass Number
• mass number (A)- the number of p+ & no in
the nucleus of an atom.
# of neutrons = mass number – atomic number
Why aren’t
electrons included
when determining
the mass number
of the atom?
Isotopes
isotope- two or more atoms having the same
atomic number (same #p+) , but different
mass numbers (due to different #no).
Isotope Notation
Nuclear Notation
Hyphen Notation
Uses the elements symbol followed by a hyphen & the mass
number.
C-12
How many protons, neutrons &
electrons are there in the following?
Cl-38
35Cl-1
Br-80
32S-2
N-14
56Fe+3
Changes in the Nucleus
Nuclear Reaction- changes that occur in the
atom’s nucleus.
• Nuclear reactions can change the
composition of an atom’s nucleus
permanently.
Types of Radioactive Decay
Alpha Radiation (a)- stream of high energy alpha particles.
• Consists of 2 protons & 2 neutrons making it identical to
a He-4 nucleus.
• Alpha particles can be represented by:
a
4
2
4
2
He
+2
4
2
He
• Most alpha particles are able to travel only a few
centimeters through air and are easily stopped by
clothing etc.
Alpha Decay
239
94
Pu 
234
92
U
parent
235
92
4
2
U + He
230
90
4
2
Th + He
daughter
Types of Radioactive Decay
Beta Radiation (b) – consists of a stream of high
speed electrons. These electrons are not
electrons that are in motion around the atom’s
nucleus.
• Beta particles can be represented by:
0 -1
1
e
0
1
e
0
1
b
• Can penetrate through clothing and damage
skin.
Beta Decay
6
2
He  Li + b
24
11
6
3
0
-1
Na  Mg + b
parent
24
12
0
-1
daughter
Checking for Understanding
alpha
decay
beta
decay
210
84
Po  He +
14
6
4
2
C
14
7
206
82
Pb
N+ b
0
-1
226
88
Ra
The Mole
mole (mol)- SI Unit for
the amount of a
substance that
contains as many
particles as there are
atoms in exactly 12g
of carbon-12.
• A unit of counting, like
the dozen.
Avogadro’s Number
Avogadro’s Number - the number of
particles in exactly one mole of a pure
substance.
1 mole = 6.0221415 X 1023
1 mol = 6.02 x
23
10
Amedeo Avogadro
Atomic Mass
atomic mass - the mass of one mole of an
atom
• Atomic mass is expressed in atomic mass
units (amu) or (u) or g/mol.
• Can be found on the periodic table.
• All atomic masses are based on the
atomic mass of carbon-12 being 12 amu.
Molar Mass
molar mass - the mass of one
mole of a pure substance.
• Molar mass is written in units
of amu or g/mol.
Atomic mass vs. Molar mass
• atomic mass - the mass of one
mole of an atom.
• molar mass - the mass of one
mole of a pure substance.
Atomic Mass vs. Molar Mass
Example
Atomic Mass
Molar Mass
Na
22.99 g/mol
22.99 g/mol
107.87 g/mol
107.87 g/mol
12.01 g/mol
12.01 g/mol
16.00 g/mol
16.00 g/mol
Ag
C
O
Molar Mass of Compounds
Compound
H2O
C6H12O6
Molar Mass
18.02 g/mol
180.18 g/mol
NaCl
58.44 g/mol
Cl2
70.90 g/mol
(NH4)3PO4
149.12 g/mol
CuSO4·5H2O
249.72 g/mol
The Mole Bridge
Cheer
• I say grams, you say molar mass.
grams – molar mass
grams – molar mass
Grams to Moles
Converting grams to moles: divide by molar mass.
1. How many moles of Ca are in 5.00g of Ca?
1 mol Ca
5.00g Ca x
= 0.125 mol Ca
40.08 g Ca
Moles to Grams
Converting from moles to grams: multiply by molar mass
1. What is the mass in grams of 2.25 moles of Fe?
55.85 g Fe
2.25 mol Fe x
= 126 g Fe
1 mole Fe
Types of Particles
• Atoms – C, Cu, He
• Molecules – O2, C12H22O11, CO2 (all
nonmetals in the formula)
• Formula units – NaCl, CaCl2, Mg(NO3)2
(includes a metal in the formula)
1 mole = 6.02 x 1023 particles
Particles to Moles
Converting particles to moles: divide by
Avogadro's Number.
1. How many moles of Pb are in 1.50 X
1025 atoms of Pb?
2.49 x 101 moles Pb
Moles to Particles
Converting moles to atoms: multiply by
Avogadro's Number.
1. How many molecules of NO are in 0.87
moles of NO?
23
5.2 x 10
molecules NO
Grams to Moles to Particles
Example: How many molecules of N2
are in 57.1g of N2?
23
1 mol N 2
6.02
x
10
molecules
N
2
57.1 g N 2 x
x
28.02 g N 2
1 mol N 2
= 1.23 x 1024 molecules N 2
Particles to Moles to Grams
Example: How many grams of NaF are
in 7.89 X 1024 formula units of NaF?
24
7.89 x 10 f.un. NaF x
1 mol NaF
41.99 g NaF
x
23
6.02 x 10 f.un. NaF
1 mol NaF
= 550. g NaF
Chapter 4
Arrangement of
Electrons in Atoms
The emission of light is fundamentally related to the
behavior of electrons.
The Concept of Energy
• Energy – the ability (capacity) to do work
• Units of Energy: calorie (cal), Joule (J)
• Work is done when an object moves some
distance in response to a force (pushing or
pulling).
work = force x distance
Forms of Energy
•
•
•
•
•
•
•
•
•
Mechanical energy
Chemical energy
Solar energy
Radiant or Electromagnetic energy
Electrical energy
Nuclear energy
Thermal energy
Sound energy
Magnetic energy
Where Does Energy Go?
Law of Conservation of Energy – energy
cannot be created nor destroyed
Energy is continually transferred from one
thing to another, never disappearing.
Types of Energy
• kinetic energy – the energy that objects
have because they are moving
KE – “energy of motion”
• potential energy – the energy that is
available for doing work at some later
time
PE – “energy of position”
Energy Units
• What units are used for energy?
 joule = J (SI unit)
 calorie = cal (non SI unit)
 Calorie = Cal (food calorie =
1000 cal)
1.00 cal = 4.184 J
Example: How many calories of energy is 1067 joules?
1.00 calorie
1067 J x
 255. calories
4.184 J
James Prescott Joule
Measurable Properties of Light
• Wavelength (l) – the distance between
corresponding points on adjacent waves
Units: meters, nanometers, etc.
1m = 1x109nm
Measurable Properties of Light
• Frequency (u)- number of waves that
pass a given point in a specific time
(usually one second)
-1
Unit: Hertz (Hz)
1
Hz = or s
s
Types of Electromagnetic Radiation
•
•
•
•
•
•
•
Gamma rays
X-rays
Ultraviolet
Visible
Infrared
Microwaves
Radio waves
High energy, E
High frequency, u
•Violet
•Indigo
•Blue
•Green
•Yellow
•Orange
•Red
High wavelength, l
Memorize these in order. Know E, u, and l order.
Relationship between u and l
c = lu
c = speed of light in m/s
l = wavelength in m
u = frequency in 1/s
What is the constant, h?
Planck's constant, h = 6.626 x 10
E  hu
E = energy in J
h = constant in J.s
u = frequency in 1/s or Hz
-34
J s
Combining the Equations you
get…
c =l u
l
E = hu
Check for Understanding
13. What would the energy be for light with
a frequency of 5.68 x 1012 1/s?
14. What is the energy of light that has a
wavelength of 7.89 x 10-11 m?
Check for Understanding
13. E = hv
1
E = (6.626 x 10 J  s)(5.68 x 10 )
s
E = 3.76 x 10-21J
-34
14. E =
12
hc
l
m
(6.626 x 10 J  s)(3.00 x 10
)
s
E=
-11
7.89 x 10 m
-15
E = 2.52 x 10 J
-34
8
Photoelectric Effect
• Photoelectric effect- the emission of electrons
from a metal when light shines on the metal.
• The photoelectric effect does not occur when
the light’s frequency is below a certain amount
regardless of the intensity of the light.
Absorption vs. Emission
• Absorption – when energy is “taken in” by
electrons
• Emission – when energy is “given off” by
electrons
Energy States of an Electron
• Ground State- the lowest
energy state of an atom.
(stable)
• Excited state- state in
which an atom’s potential
energy is increased from
that of the ground state.
(unstable).
Spectroscopy
• Spectroscope – used to
separate light into a spectrum
by wavelength so it can be
examined.
• Line-Emission Spectrum –
produced when an electron
jumps from a higher energy
level to a lower energy level.
Acts as an atomic fingerprint.
How to use a Spectroscope
Discovery of the
Line-Emission Spectrum
• Hydrogen atoms were excited by passing
a high-voltage current through hydrogen
gas causing the gas to glow a lavender
color.
• When viewed with a spectroscope
(diffraction grating or prism) the lavender
light separated into four narrow lines of
different color.
Hydrogen Line Spectrum
• Each of the lines seen in the hydrogen
spectrum is a result of light at a different
wavelength.
• Since light of a particular wavelength has
a definite frequency and a definite energy,
the lines of the hydrogen spectrum must
be a result of the emission of photons with
specific energies.
Other Emission Spectrum
helium
carbon
The Bohr Model
Bohr Model of the Atom
• The fact that hydrogen atoms only
released photons of specific frequencies
indicated that differences between the
atom’s energy levels were fixed.
• The puzzle behind the hydrogen-atom
spectra was solved in 1913 by Niels Bohr.
Bohr Model of the Atom
• The orbits are separated from
one another by empty space
where the electrons cannot exist.
• The energy of the electrons
becomes higher as they get
farther away from the nucleus.
• Electrons can move from one
energy level to the next by
gaining or losing a finite amount
of energy.
Chapter 4 – Part II
Arrangement of
Electrons in Atoms
The emission of light is fundamentally related to the
behavior of electrons.
Heisenberg Uncertainty
Principle
• In 1927, Werner Heisenberg stated that it is
impossible to simultaneously determine both the
position and velocity of an electron or any other
particle. This became known as the Heisenberg
Uncertainty Principle.
• Quantum theory was more widely accepted after
this proposal.
Quantum Model of the Atom
The dots represent 1 electron and the
region in which you are likely to find it
90% of the time.
Energy Levels
• Quantum numbers describe the location
of electrons in an atom.
Principal Quantum Number, n
• principal quantum number, n – indicates
the main energy level occupied by the
electron
• n = 1→7
(notice how many periods are
on the periodic table…7 !)
(Does this sound like the Bohr model?)
Energy Sublevels
No more
Bohr model
• Every energy level can be broken down
into sublevels.
• Each sublevel has an energy that is
slightly different.
• Not the traditional step ladder anymore,
but still you can’t have energies in
between the rungs of our new ladder.
Orbitals
• orbital – a 3-D region in space within an
energy level in which the electron is most
likely to be found (90% of the time).
An orbital can hold a
maximum of 2 e- !
Pauli Exclusion Principle
Sublevels in the Quantum
Model
s orbital
2 electrons
1 orientation
Sub Levels in the Quantum
Model
p orbitals: px, py, pz
6 electrons total
3 orientations
Sub Levels in the Quantum
Model
d orbitals: dxy, dxy, dyz, dx2-y2, dz2
10 electrons total
5 orientations
Sub Levels in the Quantum
Model
f orbitals
14 electrons total
7 orientations
Summary of orbitals
energy
levels
n=1
n=2
n=3
n=4
n=5
n=6
n=7
sublevels
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
5s, 5p, 5d, 5f
6s, 6p, 6d
7s, 7p
Orbital
Hund’s Rule
– within
equal
energy
orbitals, the
e- are
distributed
to have the
maxiumum
unpaired epossible
Pauli Exclusion Principle –
Aufbau
AufbauPrinciple
Principle
a maximum of 2 e- may
– –fillfillininthe
thelowest
lowest
Diagrams
occupy one orbital, both
possible
possibleenergy
energy
with opposites spins
orbital
orbital
3d
4s
3s
2s
1s
Pauli Exclusion Principle –
Aufbau Principle
3p
a maximum of 2 e- may
– fill in the lowest
occupy one orbital, both
possible energy
with opposite spins
orbital
2p
Valence Eletrons
valence electrons – electrons in the highest
energy level
Br: 1s22s22p63s23p64s23d104p5
1st E level
2nd E level
3rd E level
4th E level
4th is the highest Energy level,
so there are 7 valence electrons
Chapter 5:
Periodic Trends
Who told the
elements where to
go?
MENDELEEV!
(Father of the periodic table)
History of the Periodic Table
Early 1800’s
J.W. Dobereiner organized
elements into triads.
Triads- groups of three
elements with similar
properties. The middle
element of the triad was
thought to be an approximate
average of the properties of
the first & third element.
History of the Periodic Table
Mid 1800’s
J.A.R. Newlands- developed the law of
octaves in 1865.
Law of octaves- stated that if elements
were arranged by increasing atomic
mass, the properties of the eighth
element were similar to properties of
the first, the ninth like those of the
second, the tenth like those of the
third.
History of the Periodic Table
Mid 1800’s
In 1869, Dmitri Mendeleev published
the first periodic table.
• Mendeleev arranged the elements
horizontally by increasing atomic
mass and placed elements in
groups (vertically) based on
similar properties.
History of the Periodic Table
Early 1900’s
In 1913, Henry Moseley
developed the Modern
Periodic Table. He
determined the atomic
numbers of many elements
and then used those values to
arrange the periodic table in
rows by increasing atomic
number and in columns by
similar properties.
Metals, Nonmetals, & Metalloids
•
•
•
•
•
•
Metals
Make up most of the periodic table.
Ductile- can be drawn into wire
Malleable- can be hammered into thin sheets.
Lustrous- shiny
Good Conductors of heat & electricity
Located to the left of the step ladder on the
periodic table.
Metals, Nonmetals, & Metalloids
•
•
•
•
Nonmetals
Brittle-break when hammered.
Poor conductors of heat & electricity.
Lack luster
Located to the right of the step ladder on the
periodic table.
Metals, Nonmetals, & Metalloids
Metalloids
• Semimetals.
• Properties of both metals & nonmetals.
• Located along the step ladder on the
periodic table.
• Examples: B, Si, Ge, As, Sb, Te
Metals, Nonmetals, & Metalloids
Alkali Metals- Group 1
• Most reactive group of metals.
• Usually found in combined
form as a salt due to their high
reactivity.
• Combine vigorously with
nonmetals especially groups
16 & 17.
• React readily with water.
• Soft and silvery appearance.
Alkaline Earth Metals- Group 2
• Found in the earth’s crust but not
in the elemental form due to their
high reactivity. They are usually
found in rock structures.
• 2nd most reactive group of
metals.
• More dense than group 1.
• Shiny silvery-white color.
Transition Metals- Groups 3-12
• Compose the d-block.
• Typical metallic
properties.
• Good conductors.
• Lustrous.
• Produce colored ions.
Main Group Elements Groups 13-18
• Compose the p-block.
• Properties of elements
vary greatly.
• Contains all of the
nonmetals & metalloids as
well as some metals.
Halogens -Group 17
• Most reactive groups of non-metals.
• React vigorously with metals (especially
groups 1 & 2) to produce salts.
• Fluorine is a poisonous pale yellow gas,
chlorine is a poisonous pale green gas,
bromine is a toxic and caustic brown volatile
liquid, and iodine is a shiny black solid which
easily sublimes to form a violet vapor on
heating.
• Found in nature in the combined state.
Noble Gases- Group 18
•
•
•
•
Least reactive of all elements.
Often called inert gases.
All are gases.
The noble gases are all found in minute
quantities in the atmosphere, and are
isolated by fractional distillation of liquid air.
Helium can be obtained from natural gas
wells where it has accumulated as a result
of radioactive decay.
Inner Transition Metals- Periods 6 & 7
• Compose the f-block.
• Fill in Between Groups 3 & 4 on the
Periodic Table.
Lanthanides (Period 6)- Rare Earth
Metals
• Shiny reactive metals
Actinides (Period 7)
• Unstable & radioactive metals.
• Most are laboratory made.
Reactivity of Metals Trend
Period Trend- Metals increase in
activity from right to left on the
periodic table.
• The alkali metals are the most
reactive group of metals.
Group Trend- Metals increase in
reactivity from top to bottom with
a group.
• Ra is the most reactive alkaline
earth metal.
Reactivity of Nonmetals Trend
Period Trend- Nonmetals increase
in activity from left to right on the
periodic table with the exception
of the noble gases.
• The halogens are the most
reactive group of nonmetals.
Group Trend- Nonmetals increase
in reactivity from bottom to top
with a group.
• F is the most reactive halogen.
Reactivity Trends
What is the most reactive metal on the
periodic table? Explain.
Circle the most reactive nonmetal in each
row:
1. Te
Po S
2. Br
I
Cl
Atomic Radii
Atomic Radius- ½ the distance between the nuclei of
two identical atoms that are bonded together.
• Measure from nucleus to nucleus and divide by two.
• The larger the atomic radius the larger the atom.
Atomic Radii Trend
Atomic Radii increases as you travel down &
left on the periodic table!
Atomic Radii Trend
Atomic Radii Trends
1. Circle the atom with the largest atomic radii.
a. I F
Cl
b. O
B
N
c. S
Sb
Sr
2. Put the following elements in order of increasing atomic
radii.
a. Na, P, Rb
3. Which element has the largest atomic radii?
Ionic Radii
Cation- a positive ion formed from the loss of
electrons. Ex. Ca+2
• Metals typically form cations.
• Cations are smaller than their parent atoms.
Example: Na > Na+
Ionic Radii
Anion- a negative ion formed from the gain of
electrons. Ex. Cl-1
• Nonmetals typically form anions.
• Anions are larger than their parent atoms.
Example: Cl-1 > Cl
Ionic Radii Trends
1. Circle the larger ion.
a. Cl-1
Cl
b. Na
Na+1
c. N
N-3
d. Fe+2
Fe+3
e. C+4
C-4
Ionization Energy
Ionization Energy- the energy required to remove an
electron from an atom.
• Nonmetals have higher ionization energies than
metals.
• The lower the ionization energy of an atom the easier
it is to remove an electron.
Ionization Energy Trend
General Group Trend (representative elements)Ionization energy increases as you travel up a
group in the periodic table.
Reason: As you travel up a group there are less
principle energy levels, therefore, electrons are
located closer to the nucleus. It is harder to
remove electrons that are closer to the nucleus
to due the nuclear charge attraction.
Ionization Energy Trend
Ionization energy increases as you travel up &
right on the periodic table!
Ionization Energy Trends
1. Circle the atom with the highest ionization energy.
a. F
O
B
b. As
N
P
c. F
C
Li
Electron Affinity
Electron Affinity- energy change that occurs when an
electron is acquired by a neutral atom.
• Nonmetals have a more negative (LOWER) electron
affinity than metals.
• Metals typically have a more positive (HIGHER)
electron affinity than nonmetals.
• Noble gases have an electron affinity between metals
& nonmetals.
Electron Affinity Trends
1. Circle the atom with the highest electron affinity.
a. N
S
Ar
b. O
Ne
Ca
c. Cu
Ar
Kr
Electronegativity
Electronegativity- measure of the ability of an atom in a
chemical compound to attract electrons.
• The higher the electronegativity the greater the
attraction for electrons.
• The electronegativity scale was determined by Linus
Pauling.
• The electronegativity scale ranges from 0 to 4.
• Since some noble gases do not form compounds
they have electronegativity values of 0.
• Used to determine the polarity of an bond.
Electronegativity Trend
Electronegativity increases as you travel up &
right on the periodic table!
Electronegativity Trends
1. Circle the atom with the highest electronegativity.
a. F
S
Sn
b. Li
N
O
Chemical Bonding
Chapter 6
Pages 174-213
15-1 The Attachment Between
Atoms
atoms combine to form
ionic bonds
(M + NM)
covalent bonds
(NM + NM)
chemical bond – a mutual electrical
attraction between the nuclei and valence
electrons of two atoms that binds the
atoms together
Ionic Bonding
• ionic bond – when electrons are taken by one
atom from another atom
metal and a nonmetal
NaCl
cation and anion
(The charges are “hidden” to make a neutral compound.)
The simplest ratio of the packed
ions is called:
The Formula Unit
Ex: NaCl
Ions
• cations (+)
• anions (-)
• monatomic ions – ions formed from one
atom
Examples: Na+ or O-2
• polyatomic ions - ions formed from two or
more atoms bonded together
Examples: NH4+ or SO4-2
Ionic Compounds
• solid at room temperature
• high melting points (thus are usually solid at RT)
• formula unit represents the lowest ratio of ions
that combine to form a neutral compound
• most are crystalline solids
• when dissolved in water, the ionic compounds
will break up into ions (dissociate)
• the solutions of ionic compounds will conduct
electricity (electrolytes)
Covalent Bonding
covalent bond – when electrons are shared
between two atoms
– two nonmetals
– No ions formed! (no electrons are taken)
There is
another type
of bond, not
purely
covalent
and not
purely ionic.
Nonpolar
Pure
Covalent
Polar
Covalent
Ionic
Electronegativity
• electronegativity – a measure of the ability
of an atom in a chemical compound to
attract electrons from another atom in the
compound
The difference in electronegativity values for
two atoms will indicate whether the two
atoms form an ionic bond or a polar or
nonpolar covalent bond.
Electronegativity Differences
•
•
•
0.0 to 0.3 nonpolar covalent
0.3 to 1.7 polar covalent
1.7 and up ionic
Learn these values…Although the general rule is
a metal and nonmetal will form an ionic bond
and two nonmetals will form a covalent bond.
Ionic, Polar Covalent, or
Nonpolar Covalent?
What kind of bond would each pair form?
1. N and S
2. S and C
3. Mg and Cl
4. C and F
5. Ba and O
VSEPR Theory
Valence
Shell
Electron
Pair
Repulsion
Repulsion between the sets of
valence-level electrons
surrounding an atom causes
these sets to be oriented as far
apart as possible.
Theory
Regions of Electron Density
What is a Region of electron density?
• Single bond (2e- connecting 2 atoms)
• Double bond(4e- connecting 2 atoms)
• Triple bond(6e- connecting 2 atoms)
• Lone pair (unbonded pair) (2e- alone on an
atom)
LINEAR
180o
sp hybridization
2 Regions of Electron Density
2 Bonds
bonded pair
of electrons
bonded pair
of electrons
TRIGONAL PLANAR
120o
sp2 hybridization
3 Regions of Electron Density
3 Bonds
3 bonded pairs
of electrons
TETRAHEDRAL
109.5o
sp3 hybridization
4 Regions of Electron Density
4 Bonds
4 bonded pairs
of electrons
TRIGONAL PYRAMIDAL
107o
sp3 hybridization
4 Regions of Electron Density
3 Bonds & 1 Lone Pair
1 lone pair
of electrons
3 bonded pairs
of electrons
BENT
105o
sp3 hybridization
4 Regions of Electron Density
2 Bonds & 2 Lone Pairs
All of these have 4 regions of electron
density and sp3 hybridization
(although the number of bonded pairs is different)
TRIGONAL BIPYRAMIDAL
120o & 90o
5 Regions of Electron Density
5 Bonds
OCTAHEDRAL
90o
6 Regions of Electron Density
6 Bonds
SF6
SQUARE PLANAR
90o
6 Regions of Electron Density
4 Bonds & 2 Lone Pairs
ICl4-
Hybridization
Regions of e- Density
2
3
4
Hybridization
sp
sp2
sp3
Chapter 6: Chemical Bonding
Part II – Polarity
What do the differences in
electronegativity indicate between
two atoms?
the bond type that they will form
What are the three types of bonds?
–Ionic (difference of 1.7+)
–Polar covalent (difference of 0.3-1.7)
–Nonpolar covalent (difference of 0.0-0.3)
Nonpolar Covalent Bond
• Nonpolar covalent bond – a covalent bond
in which the electrons are shared equally
Example: Cl2
Both chlorines
have the same
electronegativity.
Polar Covalent Bond
• Polar covalent bond – a covalent bond in
which the electrons are shared unequally
Example: HCl
H
Cl
Which end represents the hydrogen end
& which end represents the chlorine end?
How do you know this answer?
Cl has a higher
electronegativity
Dipole
• dipole – partial negative or partial positive
charge formed during unequal sharing of
electrons (in polar bonds only)
• d+
d• The direction of a dipole is from the
dipole’s positive pole to its negative pole.
Nonpolar vs. Polar Covalent Bond
There is no
dipole drawn
over the H2
molecule…
WHY?
There is not an
uneven
distribution of
electrons.
The dipole is
drawn over the
HCl molecule.
The arrow
points towards
the more
electronegative
element, Cl.
Nonpolar vs. Polar Covalent Bond
Compare the
electron
clouds for the
H2 and HCl
molecules.
Why are they
different?
The electrons
are not shared
equally within
the HCl
molecule.
Do we usually
draw the
electron
clouds with the
Lewis
NO
structures?
Dipole
Example: PCl3 has three
dipoles
These
threebecause
dipoles are
there are three polardrawn
bonds.
beside the bonds,
What is the more
electronegative
element, P or Cl?
So, how would the
dipoles be drawn?
pointing towards the
more electronegative
element, Cl.
P
Cl
Cl
Cl
Each dipole represents a polar covalent bond.
You Try It!
•
Draw the Lewis Structures on your white
board for the following molecules
(Aha!…they must all be covalent, then!)
and LABEL ALL THE DIPOLES along
each bond.
Draw them on your white
1. CO2
board and only when you are
2. HBr
ready to check the Lewis
3. NH3
structures AND dipoles, go on
to the next slide.
Answers
Did you make CO2
linear?
Did you make NH3
trigonal pyramidal?
Did you show the
dipoles pointing to the
more electronegative
element?
Dipoles within Molecules
• There are dipoles created along bonds,
however, a molecule can exhibit overall
polarity.
Here is the
dipole along the
first H-O
bond…
Here is the
dipole along the
second H-O
bond…
An overall
dipole is
created.
Two individual dipoles along each bond cause the
molecule to have a greater electron cloud towards
the oxygen end leaving the other end of the
molecule more positive. This causes a partial
negative end towards the oxygen and the a partial
positive end towards the hydrogens (you can only
have two “ends” to a small molecule.
An overall
dipole is
created.
When you have a
two atom
molecule, it is
The bond
easy to tell if it is a is polar
polar molecule.
and so is
the overall
molecule.
partially negative end
partially positive end
Polarity of a Molecule
• Polarity of a molecule depends upon two
things
– the polarity of the bonds
– the shape of the molecule
• All molecules with only nonpolar bonds are
nonpolar molecules.
• Molecules with polar bonds, may or may
not be polar molecules, depending on the
shape.
carbon monoxide, CO
Is the bond polar?
YES
Is the molecule polar? YES…a two atom
molecule is polar if
the bond is polar.
ammonia, NH3
Are the bonds polar?
YES
YES…the molecule is
not symmetrical and
does have a partial
positive end and a
partial negative end.
You can draw on
Is the molecule polar?
overall dipole.
methane, CH4
Are the bonds polar?
Is the molecule polar?
YES
NO…the molecule is
symmetrical and
does NOT have a
partial positive end
and a partial
negative end. (The
center and the
outside does not
count as an “end”.)
Polar or Nonpolar Molecule?
Rule #1: All molecules with nonpolar bonds
are always nonpolar molecules.
Rule #2: Molecules with polar bonds, will be
polar only if they are nonsymmetrical (like
NH3).
You Try It! #2
Draw the following molecules, including the
dipoles along each bond.
Determine whether they are polar or
nonpolar molecules.
• CH3Cl
Draw them on your
white board. Click to
• SF6
see the answers.
• O
2
• H2S
• KCl
Chapter 7 – Part I
Formula Writing and Naming
May I please have a glass of
Be sure to know how to name…
• Ions
– polyatomic and monotomic, cations and anions
• Ionic compounds
– ones with and without Roman numerals
• Molecular/Covalent compounds
– use prefixes for these (mon-, di-, tri-, tetra-, …)
• Acids
Chapter 7 – Part II
Math with Formulas
Percent Composition
• What is the percent composition of oxygen
in carbon dioxide?
Percent composition is determined by mass. You need to know the
total mass of the element for the bottom of your fraction and the mass
of the part of the compound in question for the top of the fraction.
multiplied by 2 because there
are 2 oxygen in the formula
molar mass of oxygen
2 x 16.00 g
%O =
x 100.% = 72.7% O
44.01 g
molar mass of carbon dioxide
Empirical Formulas
• What is the empirical formula for glucose,
C6H12O6?
An empirical formula is just the lowest ratio of the molecular
formula. It is not representative of the real molecule, but it just
shows the ratio of atoms in that molecule.
C6H12O6 has a ratio of 6:12:6 which can be
reduced, so the answer is
empirical formula is CH2O (1:2:1 ratio)
Determining Molecular Formulas
from Empirical Formulas
• What is the molecular formula for a compound
that weighs 110.g/mol and has an empirical
formula of C3H3O?
The molecular formula must weigh some multiple of the
empirical formula. First determine the mass of the empirical
formula.
C3H3O weighs 55 g/mol and the real compound
weighs 110. g/mol, or twice that much.
So, the real formula is 2x C3H3O or
C6H6O2