Download Chemistry Ch3 Honors

Atoms:
The Building Blocks of Matter
Lab
• Boxes of the Unknown
Do Now
• What is a theory?
• What is a model?
• How do you make inferences about things you
can not see?
Atom
• From Greek, meaning
“indivisible”
• Atomic theory (from
400 BCE) = atoms are
the building blocks of
matter
• There wasn’t any
evidence for nearly
2000 years.
Atomic Size
• At sea level, one cubic
centimeter of air (size of a
sugar cube, or marble) will
have 45 billion billion atoms
within it.
– 45,000,000,000,000,000,000
• How many atoms would it take
to fill a universe?
• If you tried to count to
45,000,000,000,000,000,000 it
would take you 400,000 years
• Fill Rye COMPLETELY with
45,000,000,000,000,000,000
marbles.
Atomic Size
• To see the atoms in a drop
of water, you would need
to enlarge the drop until…
it is 24 kilometers wide!
• Think of a line 1 millimeter
long. If this line were
blown up to the size of the
empire state building, an
atom would be…
a tenth the thickness of a
sheet of paper.
Some History
• Democritus
– 460-371 B.C.
– ancient Greek philosopher
– believed all matter consisted of
extremely small particles that
could not be divided
– atoms, from Greek word
atomos, means “uncut” or
“indivisible”
• Aristotle
– believed all matter came from
only four elements—earth, air,
fire and water
Some Scientific Laws
• The Law of Definite Proportions
• The Law of Conservation of Mass
• The Law of Multiple Proportions
Law of Definite Proportions
• Two samples of a chemical compound contain
the same elements in exactly the same
proportions by mass regardless of the size of
the sample or source of the compound.
– NaCl = 39.3% sodium 60.7% chlorine
– H2O = 11.2% hydrogen 88.8% oxygen
– C2H6O2 = 38.7% carbon, 9.7% hydrogen, 51.6% oxygen
Law of Conservation of Mass
• Mass is neither created nor destroyed during
ordinary chemical reactions or physical
changes
• Thus, the mass of the reactants equals the
mass of the products
Law of Multiple Proportions
• If two or more different compounds are
composed of the same two elements, then
the ratio of the masses of the second element
combined with a certain mass of the first
element is always a ratio of small whole
numbers.
– NO and NO2 = 1:2
– H2O and H2O2 = 1:2
– SO2 and SO3 = 2:3
Dalton’s Atomic Theory
• Dalton proposed an explanation for the law of
conservation of mass, the law of definite
proportion, and the law of multiple
proportions.
• He reasoned that elements are composed of
one kind of atom and that only whole
numbers of two or more kinds of atoms can
combine to form compounds.
• He proposed the solid sphere model
Dalton’s Atomic Theory
• All matter is composed of extremely small particles
called atoms, which cannot be subdivided, created, or
destroyed.
• Atoms of a given element are identical in physical
(size/mass) and chemical properties.
• Atoms of different elements differ in physical
(size/mass) and chemical properties.
• Atoms of different elements combine in simple, wholenumber ratios to form chemical compounds.
• In chemical reactions, atoms are combined, separated,
or rearranged, but never created, destroyed, or
changed.
Modern Atomic Theory
• Atoms are divisible into even smaller particles.
These smaller parts of the atom are called
subatomic particles
– Electrons
– Protons
– Neutrons
Discovering Electrons
• In 1897, J.J. Thomson used a cathode ray tube
(passes electricity through a glass tube with
little pressure) to deduce the presence of a
negatively charged particle.
J.J. Thomson’s cathode ray tube
• He knew that rays must have come from the
atoms of the cathode because most of the
atoms in the air had been pumped out of the
tube. Because the cathode ray came from the
negatively charged cathode, Thompson
reasoned that the ray was negatively charged.
J.J. Thomson’s cathode ray tube
• He observed that when a small paddle wheel was placed in
the path of the rays, the wheel would turn. This observation
suggested that the cathode ray consisted on tiny particles that
were hitting the paddles of the wheel.
• His experiments showed the cathode ray consists of particles
that have mass and a negative charge. These were called
electrons.
• He proposed the plum pudding model.
http://www.bbc.co.uk/history/british/victorians/launch_ani_paddle_steamship.shtml
Discovering the Nucleus
• In the 1900s, Ernest Rutherford performed his
gold foil experiments
– He directed small, positively charged alpha
particles (that are helium nuclei) at a thin gold foil.
– Particle hits on the detecting screen (film) were
recorded and deflected angles were measured.
Rutherford’s gold foil experiments
This diagram shows the expected
result of Rutherford's experiment if
the "plum pudding" model of the
atom is correct.
This diagram shows the actual
result. Most of the alpha particles
are only slightly deflected, as
expected, but occasionally one is
deflected back towards the source
Only a very concentrated (dense) positive charge in a tiny space within the gold
atom could possibly repel the fast-moving, positively charged alpha particles
enough to reverse the direction of the alpha particles.
Rutherford’s gold foil experiments
• His experiments showed that the nucleus is very
small and positively charged.
• He also hypothesized that the mass of the
nucleus must be larger than the mass of the
alpha particles, otherwise the alpha particles
would have knocked the nucleus out of the way.
• He also argued that most of the alpha particles
were not deflected, because most of the atom
was empty space.
• He proposed a planetary model or nuclear model
Try it Yourself!
• In the following pictures, there is a target hidden by a
cloud. To figure out the shape of the target, we shot
some beams into the cloud and recorded where the
beams came out. Can you figure out the shape of the
target?
RUTHERFORD ACTIVITY
HALLWAY
Pennies
Rolled Marbles
Rolled Marbles
http://www.learner.org/resources/series61.html
Pennies
Equal
Distance
From each
other
Do Now
• What is an atom?
• What particles make up an atom?
• Where are the particles located?
The Nucleus
• Using measurements from
Rutherford’s experiment,
scientists calculated the
radius of the nucleus to be
less than 1/10,000 of the
atom.
– If the nucleus were the size of
a marble, the atom would be
the size of a football stadium
The Nucleus
• Protons = positively charged particles
– The charge of a proton was calculated to be equal but
opposite to the charge of an electron
– The mass of a proton is almost 2000x the mass of an
electron
• Neutrons = neutral particles
– The mass of a neuron is almost equal to the mass of a
proton
• The sum total of masses of protons, neutrons,
and electrons equals the mass of the atom.
Mass of atoms are measured in Atomic Mass Units!
1 amu = 1/12 mass the Carbon-12
(amu)
Atoms
• All living things are made up of tiny units
called ATOMS.
ATOMS consist of electrons orbiting around a
nucleus.
ELECTRONS
• (-) negative electrical charge found in the
space around the nucleus
NUCLEUS
• PROTON
(+) has a positive electrical charge.
• NEUTRON
has a neutral charge (no charge)
Subatomic Particles
ATOM
NUCLEUS
ELECTRONS
PROTONS
NEUTRONS
POSITIVE
CHARGE
NEUTRAL
CHARGE
NEGATIVE CHARGE
Atomic Identity
• If the number of electrons equals the number
of protons, the atom is electrically neutral.
(No electrical charge)
• Elements differ in their number of protons
and therefore in the amount of positive
charge their nuclei possess.
• The number of protons determines an atom’s
identity.
Atomic Number
Atomic # = p+
• Atomic Number (Z)= number of protons of
each atom of that element
– Atoms of different elements have different
numbers of protons (different atomic numbers)
– Atoms of the same element all have the same
number of protons (same atomic numbers)
• The atomic number identifies the element
– 113 elements have been identified, with 113
different atomic numbers
Atomic Number
• Because atoms are neutral, they must have
the equal numbers of protons and electrons.
• Therefore, the atomic number tells us how
many protons and also how many electrons an
atom has.
How many PROTONS
and ELECTRONS are in:
•
•
•
•
•
•
•
Silver
Hydrogen
Neon
Gold
Boron
Sodium
Tungsten
47
1
10
79
5
11
74
Mass Number
Mass # = p+ + n0
• Mass number = the total number of protons and
neutrons (total number of particles in the nucleus)
• Mass numbers can vary among atoms of a single
element, because atoms of the same element can have
different numbers of neutrons.
• Different elements can have the same mass numbers,
because the mass number does not help you identify
the element, the atomic number does!
Try this:
Mass # = p+ + n0
Element
p+
Oxygen
33
Phosphorus
n0
e- Mass #
What did you learn today?
What did you learn today?
• Each atom has a nucleus, with an overall positive charge,
surrounded by one or more negatively charged electrons.
• Subatomic particles contained in the nucleus include protons
and neutrons.
• The proton is positively charged, and the neutron has no
charge. The electron is negatively charged.
• Protons and electrons have equal but opposite charges. The
number of protons equals the number of electrons in an
atom.
• The mass of each proton and each neutron is approximately
equal to one atomic mass unit. An electron is much less
massive than a proton or a neutron.
Do Now
• Create an Atom activity
Nuclear Symbols
•
•
235U
92
235 is the mass number of Uranium
92 is the atomic number of Uranium
• A uranium nucleus has 92 protons.
• It also has a total of 235 neutrons and protons in its
nucleus (mass number).
• How many neutrons in an atom of Uranium-235?
• Mass # – Atomic # = # of Neutrons
• 235 (protons + neutrons) – 92 protons =
143 neutrons
FIND THE NUMBER OF NEUTRONS:
•
•
•
•
•
•
•
Sodium
Calcium
Nitrogen
Iron
Argon
Lithium
What does this tell you?
12
6
C
Modern Atomic Theory
• Atoms of a particular element do share the same
atomic number (number of protons) and identical
chemical properties but the atoms of a given
element may differ in their mass numbers (number
of protons and neutrons).
• Elements occur in nature as mixtures of isotopes.
Isotopes
• Isotopes = atoms of the same element with
different numbers of neutrons and mass
numbers
 Nuclear symbol:
Mass #
Atomic #
12
6
 Hyphen notation: carbon-12
C
Isotopes
• Isotopes = atoms of the same element with
different numbers of neutrons and mass
numbers
 Nuclear symbol:
Mass #
Atomic #
14
6
 Hyphen notation: carbon-14
C
Isotopes
© Addison-Wesley Publishing Company, Inc.
Try to determine information
about these isotopes:
• Chlorine-37
– atomic #:
17
– mass #:
37
– # of protons:
17
– # of electrons:
17
– # of neutrons:
20
37
17
Cl
Isotopes of Hydrogen
Isotope
Hydrogen–1
(protium)
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium)
Protons Electrons Neutrons Nucleus
Using Mass Numbers
• How many protons, neutrons, and electrons make up an atom of Br-80?
•
•
•
•
Protons + Neutrons = 80
Protons = 35
Electrons = 35
Neutrons = 80 – 35 = 45
• How many protons, neutrons, and electrons make up an atom of C-14?
•
•
•
•
Protons + Neutrons = 14
Protons = 6
Electrons = 6
Neutrons = 14 – 6 = 8
Ions
Are created when an atom loses or gains one
or more electrons; it acquires a charge
http://web.visionlearning.com/custom/chemistry/animations/CHE1.3-an-ions.shtml
Charge of Ion = number of protons – number of electrons
More electrons than protons = negative charge (anion)
More protons than electrons = positive charge (cation)
12
6
C +1
# of protons 
# of electrons 
Total charge 
PRACTICE IONS
Ion
Li
+1
Ni
+2
Pb
+2
Ca
+2
Cs
+1
# protons
# neutrons # electrons
Chemical
Symbol
Number of
protons
Number of
electrons
Number of
neutrons
35
36
45
Atom or Ion?
I
11
12
atom
55
78
atom
14
12
Zr
12
50
Br
53
atom
44
atom
Ce
ion
27
25
32
84
80
125
73
68
108
50
71
Sc
Pb
Ni
atom
What did you learn today?
What did you learn today?
• Atoms of an element that contain the same number of protons but a
different number of neutrons are called isotopes of that element.
• The average atomic mass of an element is the weighted average of the
masses of its naturally occurring isotopes.
Lab
• Atoms family lab
Quiz
• Intro to Atoms
Introducing MOLES
• Mole = the amount of a substance that
contains as many particles as there are atoms
in exactly 12 grams of carbon-12.
– The mole is a counting unit, like a dozen.
– The mole relates to masses of atoms and
compounds.
• Avogadro’s number = The number of particles
in exactly one mole of a pure substance.
– 6.022 x 1023 particles
Do Now
• Recall the atomic theories and models
proposed by Dalton, Thomson and Rutherford.
Discoveries about the atom
Dalton
1.
All matter is composed
Of extremely small particles which
cannot be subdivided, created or
destroyed.
2. Atoms of a given element are
identical in physical and chemical
properties.
3. Atoms of different elements have
different physical and chemical
properties.
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,
but never created, destroyed, or
changed.
JJ Thomson
What did he
discover: Electron
His Experiment:
Cathode Ray Tube
His findings:
Electrons are
negatively charged
embedded in a
positive charge.
Rutherford
What did he
discover: The Nucleus
His experiment:
GOLD FOIL
EXPERIMENT (1900’s)
His findings:
The atom is mostly
empty space.
The nucleus is small.
The nucleus is dense.
The nucleus is
positively charged
Niels Bohr
•Electrons revolve around the nucleus
in specific orbits, or energy levels.
• An atom has energy levels. Electrons
can only exist in these energy levels,
not in between.
•When an atom is in the ground state,
the electrons exist in the energy
levels closest to the nucleus.
•GROUND STATE: the lowest energy
state of an atom; the electrons
occupy energy levels closest to the
nucleus.
•If an atom receives, energy, the
atom becomes excited and electrons
jump to higher energy levels.
•EXCITED STATE: an atom with
higher potential energy than in the
ground state because electrons have
“jumped” to a higher energy level.
Solid
Sphere
Model
Electron
Cloud
Model/
Quantum
Model
This model
suggested that
electrons could be
considered waves
confined to the
space around a
nucleus.
Electron cloudsregions where
electrons are
likely to be found
Refining Nuclear Models
• In 1913, Danish physicist, Niels Bohr, refined
Rutherford's idea by adding that the electrons were
in orbits around the nucleus. Rather like planets
orbiting the sun. With each orbit only able to contain
a set number of electrons.
• He proposed a Bohr model or Orbit model
Bohr’s
Atom
HELIUM ATOM
_______
_______
_________
+
_________
N
N
+
-
__________
The Bohr Model
• 1. Electrons revolve around the nucleus in specific
orbits (shells), or energy levels.
• 2. An atom has energy levels. Electrons can only
exist in these energy levels, not in between.
• 3. When an atom is in the ground state, the
electrons exist in the energy levels closest to the
nucleus.
• 4. If an atom receives energy, the atom becomes
excited and electrons jump to higher energy
levels.
http://www.visionlearning.com/library/flash_viewer.php?oid=1347&mid=51
Ground State
• The lowest energy state
of an atom
• Electrons in the first
energy level have the
lowest potential energy
since they are located
closest to the nucleus.
Found on the periodic table
Excited State
• An atom with higher potential energy than in the
ground state because electrons have “jumped” to a
higher energy level.
• Electrons with higher potential energy occupy orbits
farther from the nucleus. The further an electron is
from the nucleus, the greater its energy!
• Atom song:
• http://www.youtube.com/watch?v=vUzTQWn
-wfE&feature=related
• Build an atom:
• http://www.classzone.com/books/earth_scien
ce/terc/content/investigations/es0501/es0501
page03.cfm
Current Atomic Model
• Electrons act like particles (because they have a mass) and waves
(because they have certain frequencies corresponding to their
energy levels)
• Electrons are located in orbitals around the nucleus that correspond
to specific energy levels
• Electron clouds = orbitals that do not have sharp boundaries, but
shows 3D region where electrons are most probable to be found.
• Wave Mechanical (Electron Cloud) Model or Quantum Model
proposed by Louis de Broglie & Erwin Schrodinger
Valence Electrons
• Electrons that occupy the valence energy level
• Valence Electrons = found in outer most
energy level
Na 2-8-1
Cl 2-8-7
• Atoms can have a maximum of 8 valence
electrons
Electron Configuration
• Arrangement of electrons
• Each atom has a distinct electron
configuration.
• The ground state electron configuration is
found on the periodic table in the lower left
hand corner of each box.
Element
Ground State
Electron
Configuration
Ion
Ion’s Electron
Configuration
Na
Na
+
Mg
Mg
+2
Fe
Fe
+3
Al
Al
+3
Li
Li
+1
Use your
Periodic Table!
Lewis Dot Diagrams
(Electron Dot Diagrams)
• Represent the arrangement of electrons
around the nucleus.
• Electrons are the DOTS.
• Nucleus is the symbol.
• ONLY REPRESENT VALENCE ELECTRONS!!
• Fill one side first, then one on each side before
you pair electrons.
Lewis Dot Diagrams
(Electron Dot Diagrams)
Na
B
O
Mg
Cl
Ne
Si
H
N
What did you learn today?
What did you learn today?
• In the wave-mechanical model (electron cloud model), the electrons are in
orbitals, which are defined as the regions of the most probable electron
location (ground state).
• Each electron in an atom has its own distinct amount of energy.
• When an electron in an atom gains a specific amount of energy, the
electron is at a higher energy state (excited state).
• When an electron returns from a higher energy state to a lower energy
state, a specific amount of energy is emitted. This emitted energy can be
used to identify an element.
• The outermost electrons in an atom are called the valence electrons. In
general, the number of valence electrons affects the chemical properties
of an element.
Do Now
• What does a mole represent?
• How can you determine the mass of a mole?
Relating masses in grams
to numbers of atoms
• Mole = the amount of a substance that contains as many
particles as there are atoms in exactly 12 grams of carbon-12.
– The mole is a counting unit, like a dozen.
– The mole relates to masses of atoms and compounds.
• Avogadro’s number = The number of particles in exactly one
mole of a pure substance.
– 6.022 x 1023 particles
• Molar mass = the mass of one mole of a pure substance;
numerically equal to the atomic mass of the element in
atomic mass units (g/mol)
Molar mass is also known as Gram-formula mass (GFM)
number of moles = given mass (g)
From your
reference table!
gram formula mass
What is the molar mass of Li?
6.94 g/mol
What is the molar mass of Hg?
200.59 g/mol
What is the mass in grams of
3.50 mol of Cu?
3.50 mol Cu x 63.55g Cu
1 mol Cu
= 222g Cu
What is the mass in grams of
3.42 mol Ag?
What is the mass in grams of
0.876 mol Pb?
A chemist produced 11.9 g of
Aluminum. How many moles of Al
were produced?
11.9 g Al x 1 mol Al
26.98 g Al
= 0.441 mol Al
How many moles of Na are in
4.01 g Na?
How many moles of Zn are in 0.674 g
Zn?
The molar mass
of an element
contains one
mole of atoms.
4.00g He, 6.94g Li, and
200.59g Hg all contain
one mole of atoms.
How many atoms is
this?
Avogadro’s number:
6.02 x 1023 particles
(atoms)
How many moles are in 32 g of S?
How many atoms are in this sample of S?
What mass of B contains the same
number of atoms as 8.0 g of Bi?
• What is the gram formula mass (GFM) of salt
(NaCl)?
• What is the gram formula mass of Cl2?
• What is the molar mass of SO2?
• What is the molar mass of a substance if 0.25
moles of a substance has a mass of 45 grams?
Honors
• Moles to number of particles conversions
Honors
How many atoms in 2.5moles of Cu?
How many atoms are in 1.37moles of Hydrogen?
Honors
• How many moles of Ag are in 3.01 x 1023
atoms of Ag?
(3.01 x 1023 atoms Ag) x 1 mol Ag
(6.02 x 1023 atoms Ag)
= 0.500 mol Ag
Honors
• How many moles of W are in 1.89 x 103 atoms
of W?
Honors
• How many moles of Ni are in 2504 atoms of
Ni?
Honors
• What is the mass in grams of 1.34 x 104 atoms
of Sb?
Do Now
• If your first quarter grade is based 10% on
homework, 10% on labs, 10% quizzes, and
70% tests, what should your grade be if you
averaged 100 on homework, 90 on labs, 80 on
quizzes, and 70 on tests?
–
–
–
–
10% x 100 = 10
10% x 90 = 9
10% x 80 = 8
70% x 70 = 49
– Total = 76 %
Atomic Mass
• Atomic mass = the weighted average of the
masses of the existing isotopes of an element.
• Don’t get these confused!
– Mass number = the total number of protons and
neutrons that make up the nucleus of an atom.
– Atomic mass includes the masses of the protons,
neutrons and electrons of atoms and isotopes.
Weighted Average
• You have a box containing two sizes of
marbles.
• 25% of the marbles have masses of 2.00 g
each
• 75% of the marbles have masses of 3.00 g
each
• Calculate the weighted average….
Calculate the weighted average
• Assume you have 100 marbles
– 25%, or 25, have a mass of 2.00 g
– 75%, or 75, have a mass of 3.00 g
–
–
–
–
25 marbles x 2.00 g = 50 g
75 marbles x 3.00 g = 225 g
Total mass = 50 + 225 = 275 g
275g / 100 marbles = 2.75 g/ marble
– A simpler method is as follows:
25% = 25/100 = 0.25
75% = 75/100 = 0.75
(2.00g x 0.25) + (3.00g x 0.75) = 2.75g
Atomic Mass
• 1 amu = 1/12 the mass of a 12C atom
Carbon = 12.011 amu
Isotope
Symbol
Carbon-12
12C
Carbon-13
13C
Carbon-14
14C
Composition % Abundance in
of the nucleus
nature
6 protons
98.89%
6 neutrons
6 protons
1.11%
7 neutrons
6 protons
<0.01%
8 neutrons
Calculating Atomic Mass
Avg Atomic Mass = (Mass x %/100) + (Mass x %/100) +…
• STEP 1: Take the mass # (in amu) of each
element and multiply by its percent
abundance divided by 100 (%/100)
• STEP 2: Add all of these values together
Calculating Atomic Mass
• Boron exists as 2 isotopes: B-10 or B-11
• B-10
10
• B-11
11
5B
5B
% Abundance
19.78%
80.22%
Calculating Atomic Mass
Atomic Mass of Boron
• STEP 1: 10 x 19.78/100 = 1.978
11 x 80.22/100 = 8.8242
• STEP 2: 1.978 + 8.8242 = 10.8022 amu
Calculating Atomic Mass
• Calculate the Atomic Mass of Chlorine:
% Abundance
• Chlorine – 35
75.53
• Chlorine – 37
24.47
Calculating Atomic Mass
• Calculate the Atomic Mass of Silicon:
% Abundance
• Si – 28
92.21
• Si – 29
4.70
• Si – 30
3.09
Calculating Atomic Mass
• Calculate the Atomic Mass of Oxygen:
% Abundance
• O-16
99.762
• O-17
0.038
• O-18
0.200
Calculating Atomic Mass
• What is the average atomic mass of Cu which
is found in nature as
69.15% Cu-63 and 30.85% Cu-65?
(0.6915 x 63) +
(0.3085 x 65) = 63.6 amu
Lab
• Beanium lab to calculate atomic mass
Quiz
• Quiz on models and calculations
Do Now
• When do electrons jump to higher energy
levels?
• What happens to the energy when the
electrons return to the ground state?
Review
• When an electron in an atom gains a specific
amount of energy, the electron is at a higher
energy state (excited state).
• When an electron returns from a higher energy
state to a lower energy state, a specific amount of
energy is emitted. This emitted energy can be
used to identify an element.
Quantum
• Electrons can only absorb or release energy in
discrete, specific amounts.
• The amounts, or bundles of energy are called
quanta (or photons) corresponding to
differences in energy levels of the
orbitals/shells.
Electrons and Light
• An atom emits energy when the electron falls from
high energy levels to lower energy levels. This energy
is in the form of electromagnetic radiation.
• If the wavelength is in the visible light spectrum, the
energy can be seen as color.
Electromagnetic radiation
Most subatomic particles behave as PARTICLES and obey the physics of WAVES.
Electromagnetic Radiation
• Wavelength () - length
of one complete wave
• Frequency () - number
of waves that pass a
point during a certain
time period
– hertz (Hz) = 1/s
• Amplitude (A) - distance
from the origin to the
trough or crest
Wavelength and Frequency
Long wavelength  small frequency
Short wavelength  high frequency
Wavelength and Frequency
•
•
•
•
Formula: c = 
c = speed of light, m/s = 3.00 x 108 m/s
 = wavelength, m
v = Frequency, 1/s
Practice Problem
• Find the frequency (Hz = 1/s) of a photon with
a wavelength of 434 nm.
• Use the formula: c = 
• Given:
Work:
v=?
v=
 = 434 nm
=
m
c = 3.00 x 108 m/s
Planck Equation
• Max Planck (late 18th century) showed that
the energy of light is proportional to its
frequency.
• Formula: E = h 
• E = energy (J, joules)
• h = Planck’s constant (6.6262  10-34 J·s)
•  = frequency (Hz)
Practice Problem
• Find the energy of a red photon with a
frequency of 4.57  1014 Hz.
• Use the formula: E = h 
• Given:
Work:
E=?
E=
h = 6.6262  10-34 J·s
 = 4.57  1014 Hz
Understanding the Atomic Model:
Quantum (wave) Mechanical
Model of the Atom
• Three Physicists in the 1920’s wanted to
determine if electrons behaved as waves as
well as particles, like light.
• Louis de Broglie (electron has wave
properties)
• Werner Heisenberg (Uncertainty Principle)
• Erwin Schrodinger (mathematical equations
using probability, quantum numbers)
Louis de Broglie
Wave Properties of Matter (1924)
• Since light waves have a
particle behavior (as
shown by Einstein in the
photoelectric effect),
then particles could
have a wave behavior.
• de Broglie wavelength
Electron Motion Around
= h
Atom Shown as a de Broglie
Wave
mv
Photoelectric Effect
• Refers to the emission of electrons from a
metal when light shines on the metal.
• Albert Einstein (1905) used this to determine
that light exists in discrete quanta (particles)
of energy.
Werner Heisenberg
Uncertainty Principle
• Electrons are detected by
their interactions with
photons
• Any attempt to locate a
specific electron with a
photon knocks the electron
off its course.
• Therefore, it is impossible to
determine simultaneously
both the position and the
velocity of an electron or any
other particle.
The more certain
you are about
where the
electron is, the
less certain you
can be about
where it is going.
The more certain
you are about
where the
electron is going,
the less certain
you can be about
where it is.
Erwin Schrodinger
Quantum (wave) Mechanical
Model of the Atom (1926)
• He proposed a
mathematical explanation
of the dual wave particle
nature of an electron
known as Schrodinger’s
Wave Equation.
• His work began a new way
of dealing with subatomic
particles, known as
quantum mechanics.
Electromagnetic Spectrum
R
red
O
orange
Y
G.
yellow
green
B
blue
I
indigo
V
violet
Light Emission
• Each move from a particular
energy level to a lower energy
level will release light of a
specific wavelength.
• When certain elements are
excited, they give off energy of
a distinctive color as the
electron fall back down to
lower energy levels. These
colors are specific and can be
used to identify the elements
(Flame Test).
Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum
…produces a “bright line” spectrum
Spectral lines
• If high voltage is applied to hydrogen gas confined in a gas
tube, called a gas discharge tube, light is emitted. If this light
is passed through a prism, a series of bright lines of distinct
colors is produced. Bohr reasoned that these different colored
bands of light were actually quanta of corresponding energy.
These quanta were emitted as electrons of hydrogen atoms
returned from their higher levels in the excited state to their
lower levels in the ground state.
Bright Line Spectra
• Bright line spectrum =
the series of bright lines
produced when excited
electrons return to their
original energy levels
• Each element has its
own unique set of
spectral lines which can
therefore be used to
identify the elements
presence.
Identify the Elements
in the Unknown
Classify the following as ground state electron
configurations or excited state electron configurations.
Element
ground state electron configurations or
excited state electron configuration
lithium
1-2
calcium
2-8-7-3 excited
chlorine
2-8-7
ground
aluminum
2-7-4
excited
neon
2-7-1
excited
sodium
2-8-1
ground
potassium
2-8-7-2 excited
exited
Element
Excited State
Electron
Configuration
Ion
Excited Ion’s
Electron
Configuration
Na
Na
+
Mg
Mg
+2
Fe
Fe
+3
Al
Al
+3
Li
Li
+1
Lab
• Spectrum lab
Quiz
• Quiz on Electromagnetic spectrum and
wavelength frequency formula
Do Now
• https://www.youtube.com/watch?v=8ROHpZ0
A70I&feature=youtube_gdata_player
Electron Configurations
• Electrons are in principal energy levels
• Principle energy levels are divided into sublevels
• Each sublevel contains a certain number of
orbitals
• Each orbital can hold 2 electrons = Pauli exclusion
principle
• The properties of electrons can be described by
Quantum Numbers
1. Principle Energy Levels (n)
• Each orbit of an atom has a fixed radius.
• The greater the radius of an orbit (the farther from
the nucleus), the greater the energy of the electrons
in that orbit. The orbits or shells are known as
principal energy levels.
• The main energy level occupied by an electron is 1
through 7 (n ranges from 1 through 7).
• The period (row) an element is found in tells us its
energy level.
2. SUBLEVELS (l)
• Each energy level has a certain # of sublevels:
s sublevel
p sublevel
d sublevel
f sublevel
• 1st energy level has 1 sublevel (s)
• 2nd energy level has 2 sublevels (s,p)
• 3rd energy level has 3 sublevels (s,p,d)
• 4th energy level has 4 sublevels (s,p,d,f)
3. ORBITALS (m)
• Each sublevel contains one or more orbitals:
s sublevel has 1 orbital
p sublevel has 3 orbitals
d sublevel has 5 orbitals
f sublevel has 7 orbitals
Quantum Numbers
(n)
(l)
(m)
px
py
pz
• Orbitals combine to form a spherical shape
2px
2py
2s
2pz
Orbital
Shapes
p = dumbell
d = Dumbell and donut
s = spherical
4. SPIN (s)
• Each orbital can contain a maximum of 2
electrons.
• Electrons spin opposite to each other
• This is known as the Pauli Exclusion Principle
Clockwise
=+½
↑↓
Counterclockwise
=-½
Quantum Numbers
• No two electrons in an atom can have the
same 4 quantum numbers.
• Each e- has a unique “address” (4 QN’s):
1. Principle (n) corresponds to the energy level (1-7)
2. Angle Momentum (l) corresponds to the sublevel
(s,p,d,f)
3. Magnetic (ml) corresponds to the orbital (x,y,z)
4. Spin (s) corresponds to the electron (+/- ½)
Quantum numbers for the first four levels of orbitals in the hydrogen
atom
n
l
Orbital
designation
ml
# of
orbitals
1
0
1s
0
1
2
0
2s
0
1
1
2p
-1, 0, 1
3
0
3s
0
1
1
3p
-1, 0, 1
3
2
3d
-2, -1, 0, 1, 2
5
0
4s
0
1
1
4p
-1, 0, 1
3
2
4d
-2, -1, 0, 1, 2
5
3
4f
-3, -2, -1, 0, 1, 2, 3
7
3
4
SPIN
SUBLEVEL
How many orbitals
can it have?
Maximum #
electrons
s
1 orbital
2 electrons
p
3 orbitals
6 electrons
d
5 orbitals
10 electrons
f
7 orbitals
14 electrons
Principal
Energy Level
Sublevels
Number of
Orbitals per
sublevel
Total number
of electrons
per sublevel
1
1s
1
2
2
2s
2p
3s
3p
3d
4s
4p
4d
4f
1
3
1
3
5
1
3
5
7
2
6
2
6
10
2
6
10
14
3
4
Principal Energy Levels
1
Maximum # of Electrons
in Energy Level
2
2
8
3
18
4
32
ORBITAL BOX DIAGRAMS
• Another method of representing the electron
configuration is with the orbital box diagram.
• Each box can ONLY contain 2 electrons and they must
have opposite spins.
•
s
p
d
f
Each orbital is represented by a box.
1 orbital
3 orbitals
5 orbitals
7 orbitals
Electron Configuration
• Electrons fill orbitals of lower energy first = aufbau principle
• Sublevel orbitals are occupied by one electron before pairing
occurs = Hund’s rule
• Before element 18, the sublevels fill in order as follows:
1s, 2s, 2p, 3s, 3p
• At this point, they begin to fill out of “order.”
The 4s fills before the 3d. The
sublevels fill order according
to the following chart:
The configuration listed on the periodic table is the
ground state electron configuration.
Element
He
O
Na
F
Al
Mg
Br
Ground State Electron Configuration
The configuration listed on the periodic
table is the ground state configuration.
He 2
1s2
O 2-6
1s2 2s2 2p4
Na 2-8-1 1s2 2s2 2p6 3s1
F 2-7
1s2 2s2 2p5
Al 2-8-3
1s2 2s2 2p6 3s2 3p1
Mg 2-8-2 1s2 2s2 2p6 3s2
Br 2-8-18-7 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
Lewis Dot Diagrams
• Go back to Lewis Dot Diagrams: Place first 2
dots together representing s sublevel, then
every dot around until all p sublevels are
filled.
H
He
Li
Be
B
C
N
O
F
Do Now
• What is the longhand ground state electron
configuration for Bromine?
• Longhand Configuration
2
2
• S 16e1s 2s
•
•
•
•
6
2p
Core e-s
Shorthand Configuration
[Ne] 3s2 3p4
S 16 e[Ar] 4s1
K 19e2 3d10 4p5
[Ar]
4s
Br 35e-
2
3s
4
3p
Valence e-s
Stability
• Electron Configuration Exceptions:
• Copper
– EXPECT:
– ACTUALLY:
[Ar] 4s2 3d9
[Ar] 4s1 3d10
Copper gains stability with a full d-sublevel.
Stability
• Electron Configuration Exceptions
• Chromium
– EXPECT:
– ACTUALLY:
[Ar] 4s2 3d4
[Ar] 4s1 3d5
Chromium gains stability with a half-full d-sublevel.
A periodic table of partial
ground-state electron
configurations
Orbital Box Diagrams
• Oxygen 8e• Electron Configuration:
2
2
4
1s 2s 2p
• Orbital Diagram:
1s
2s
2p
Practice
• C
• S
Element
Lithium
Configuration
notation
Orbital notation
1s22s1
[He]2s1
____
1s
Beryllium
____
____
2p
____
____
2s
____
____
2p
____
[He]2s2p2
____
2s
____
____
2p
____
1s22s2p3
[He]2s2p3
____
2s
____
____
2p
____
1s22s2p4
[He]2s2p4
____
2s
____
____
2p
____
1s22s2p5
[He]2s2p5
____
1s
Neon
____
2s
1s22s2p2
____
1s
Fluorine
____
[He]2s2p1
____
1s
Oxygen
____
2p
1s22s2p1
____
1s
Nitrogen
____
[He]2s2
____
1s
Carbon
____
2s
1s22s2
____
1s
Boron
Noble gas
notation
____
2s
____
____
2p
____
1s22s2p6
[He]2s2p6
____
1s
____
2s
____
____
2p
____
Hund’s rule
ASSIGNING QUANTUM NUMBERS
• The shape, size, and energy of each orbital is
a function of 4 quantum numbers which
describe the location of an electron within an
atom or ion
• n (principal)  energy level
• l (orbital)
 shape of orbital
• ml (magnetic)  designates a suborbital
• s (spin)
 spin of the electron
(clockwise or counterclockwise: ½ or – ½)
Putting it all together
• N electron configuration 2-5
↑↓ ↑↓ ↑
↑
1s2
2p3
n=
l=
ml=
ms=
2s2
↑
( _____,______,_____,______)
Putting it all together
• N electron configuration 2-5
↑↓ ↑↓ ↑
↑
1s2
2p3
n=
l=
ml=
ms=
2s2
↑
( __2__,___1__,__0__,__+1/2_)
Quantum Numbers for Specific
Electron
Lab
• Flame lab
Flame Test Videos
• http://www.youtube.com/watch?v=NEUbBAG
w14k
• http://www.youtube.com/watch?v=jJvS4uc4T
bU
• http://www.trschools.com/staff/g/cgirtain/we
blabs/spectrolab.htm
• http://www.youtube.com/watch?v=o3nn4zqzf
6M
Test
• Test on Atomic Concepts