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Chemistry I
Electrons in Atoms
Chapter 5
Rutherford’s nuclear model did not
provide enough detail about how
electrons occupy the space around
the nucleus. In this chapter we will
learn how electrons are arranged
around the nucleus and how that
arrangement effects chemical
behavior.
In what ways did many scientists find
Rutherford’s model to be incomplete?
1.)In what ways did many scientists find
Rutherford’s model to be incomplete?
It
did not explain:
1.How the electrons were
arranged around the nucleus
2.Why they were not pulled
into the nucleus
3. The differences in chemical
behavior of the different
elements.
2.)In the early 1900’s what did scientists
observe about certain elements when
heated in a flame?
Every
element gives off
specific wavelengths of
light that can be used like
a fingerprint to identify the
element.
spectroscope

http://jersey.uoregon.edu/vlab/elements/
Elements.html
An element’s chemical behavior
is related to the arrangement of
electrons in its atoms.
Wave Nature of Light
What is electromagnetic radiation?
3.)What is electromagnetic radiation?
A
type of wave that is part
electrical and magnetic
energy acting at right angles.
Visible light is a type of
electromagnetic radiation.
What are the 4 characteristics of
waves?
4.) What are the 4 characteristics of
waves?
 1.)
wavelength – distance from
crest to crest
 2.) frequency- # of waves that
pass a given point per second
 3.) Amplitude – height of wave
from the normal resting to wave
crest
 4.) speed – the speed of light is a
constant 3.0 x 108 m/s.
5.)What unit do we use to express
frequency?
Hertz (Hz)
 The hertz unit is 1/second ( the inverse of
a second)

6.) What is the speed of light?
 The
speed of light is a constant
3.0 x 108 m/s
 The formula for the speed of
light is C = λ • v
λ
= wavelength (measured in meters)
 v = frequency (measured in 1/sec)

The unit for speed is m/s
(distance/time)
#6 continued
When
wavelength (λ)
increases then frequency (v)
decreases because the speed
of light ( c) is a constant
3.0X108m/s.

Also
When
λ decreases, V
increases
7.) List types of electromagnetic
radiation from the lowest to the
highest in energy ( page 139)
List types of electromagnetic
radiation from the lowest to the
highest in energy ( page 140)
Radio waves
 Microwaves
 Infrared
 Visible
 Ultraviolet (U.V.)
 X rays
 gamma

↓decreasing wavelength
↓increasing frequency
Particle Nature of Light
 When
an object is heated
only certain wavelengths of
light are emitted. The
instrument used to separate
light into its different
wavelengths is called a
spectroscope. The pattern of
wavelengths is called the
spectrum.

http://jersey.uoregon.edu/vlab/elements/
Elements.html
When
a substance is
heated the spectrum of
light that is given off
can be used to identify
the substance.

http://jersey.uoregon.edu/vlab/elements/
Elements.html
The
wave model of light
could not explain why
different substances emit
particular wavelengths of
light when heated
8.)What are some problems with the
wave model of the atom?
 It
doesn’t explain why only
certain wavelengths of light are
emitted when an object is
heated.
 It also does not explain the
photoelectric effect.( we will
discuss the photoelectric effect in
just a little bit)
Who is Max Planck?
9.)Who is Max Planck?
A
German physicist who,
in 1900, began searching
for the reason that only
certain wavelengths of
light are given off for
every element.
10.)What did Planck conclude?
Matter
can gain or lose
energy in only small specific
amounts called quanta.
As metal is heated, it glows. In fact, it glows
different colors as it gets hotter (white hot being
the hottest). By studying glowing metals, Max
Planck discovered that only certain wavelengths of
light are emitted at each specific temperature.
11.) What is a quantum?
 The
minimum amount of energy
that can be gained or lost by an
atom.
 The amount of energy it takes
for a electron to get from one
energy level to the next.
12.) If energy can only absorbed in
quanta, specific amounts, why
does energy appear to
continuous to us?
 b/c
quanta are extremely small.
13.) What is Planck’s equation that
demonstrates that energy is
related to the frequency of
radiation?
 Equation



E = hv
E – energy of a quantum
h- plancks constant
v- frequency
14.) What is Planck’s constant?
 6.62
x 10-34 J•s
15.) How are energy and frequency
related?
They
are directly related
As frequency increase,
energy increases.
What is the photoelectric effect?
What is the photoelectric effect?

http://www.tutorvista.com/content/physic
s/physics-iv/radiation-andmatter/photoelectric-effect-and-cell.php
16.)What is the photoelectric
effect?
When
light of a certain
frequency shines on the
surface of a metal, electrons
are ejected. Blue light always
results in electrons being
ejected. Red light of any
brightness does not eject
electrons.
17.) In 1905 how did Albert Einstein
explain the photoelectric effect?
He
proposed that all
electromagnetic radiation
(including visible light) is
both wavelike and particle
like in nature.
Continued next slide
17.) In 1905 how did Albert Einstein
explain the photoelectric effect?
He
suggested that while a
beam of light has many
wavelike characteristics, it is
also like a stream of particles
called photons.
18.) What is a photon?
A
photon is a particle of
electromagnetic radiation
with no mass that carries a
quantum of energy.
19.)How did Einstein use Planck’s
idea of quantum energy to
explain the photoelectric effect?
 Planck
proposed that E=hv.
Energy is equal to his constant
times the frequency of radiation.
 Einstein proposed that since blue
light has a larger frequency, it
has a larger quantum of energy.
 Continued next slide
19.) How did Einstein use Planck’s
idea of quantum energy to
explain the photoelectric effect?
E = hv for blue light is more
than E=hv for red light
because blue light has a
larger frequency than red
light.
Practice Problems page 143
20.)What is the atomic
emission spectrum?
The
atomic emission
spectrum of an element is the
set of frequencies of (light)
electromagnetic waves
emitted ( given off) by atoms
to that element when it is
heated.
21.) How does the atomic emission
spectrum for one element compare
to another element?
Each
is unique and can be
used to identify the element
like a fingerprint.
22.) Are atomic emission spectrums of
elements continuous or distinct
individual wavelengths of light?
 They
are distinct, individual
wavelengths of light.
23.) In a line spectrum which line has
the highest energy?
The
line farthest to the blue
end of the spectrum. The line
corresponding to the shortest
wavelength/ highest
frequency.
24.) How does the quantitization of
energy ( energy only exists in
specific sizes) help explain line
spectrum?
 The
energy of the photon of light
that is emitted is tied to a
specific frequency/color of light
 E = hv Each frequency (v)
corresponds to a specific color
Questions page 145
Section 2 Quantum Theory and the
Atom
 Scientists
concluded that
light was both wave and
particle. Scientists were able
to understand atomic
structure, electrons and
atomic emission spectra
better because of a better
understanding of light.
 Scientists
wanted to know
why atomic emission
spectrums are not continuous
but are discontinuous.
23.)Who was Niels Bohr?
A
Danish physicist who
worked in Rutherford’s
laboratory. He proposed the
quantum model of the atom
that helped explain why
atomic emission spectrums
only contain certain
frequencies of light.
25.)What did Niels Bohr propose?
He
proposed that a
hydrogen atom has only
certain allowable energy
levels.
The lowest level is called
the ground state.
continued
25.)What did Niels Bohr propose?
All
other levels are called
excited states.
Electrons move around the
nucleus in only certain
allowed orbits.
Bohr labeled these orbits.
continued
25.)What did Niels Bohr propose?
The
orbit closest to the
nucleus is labeled n=1 and
was lowest in energy.
26.) How did Niels Bohr explain line
spectrum with the quantization
of energy for electrons. ( How
did Bohr explain that certain
wavelengths of light are released
when an element is heated by
proposing that electrons can only
be at certain energy levels.)
#26
He
said that when an element
is at ground state ( the
lowest energy level) no
energy (light) is released. But
when an electron becomes
excited, it jumps up to
another energy level.
#26
When
the electron returns to
the lower energy level the
energy is released. The
energy (light) that is released
is equal to the difference
between energy levels.
#26
The
amount of energy that is
released is equal to a
particular frequency of light.
27.)Do we still accept Bohr’s
model?

Bohr was correct about his
idea of quantization of energy
and his basic explanation of
atomic emission spectrum.
However he could only
predict the atomic emission
spectrum of hydrogen.
27.)Do we still accept Bohr’s
model?
His basic model of the
model is considered to be
incorrect. Today we do not
believe that the electrons
are in orbits so that we
can predict their positions.

The Quantum Mechanical Model
of the Atom
28.)Who was Louis de Broglie?
A
French physics student
that proposed an idea that
eventually accounted for
the fixed energy levels in
Bohr’s model.
29.)What did De Broglie propose?
He
proposed that all
matter moves in waves.
The larger the mass the
smaller the wavelength.
Only very small objects
have noticeable
wavelengths.
30.)What was De Broglie’s
equation?
λ = h/ m·v
DeBroglie’s equation
predicts a large mass will
have a small wavelength (λ)
only a very small mass will
have a noticeable
wavelength.

Page 149
If electrons move in waves then only
electrons with matching wavelengths
can be accepted into an energy level
and when an electron is released to a
lower energy level particular
wavelengths of energy are released.
29.)What did Heisenberg propose?
It
is impossible to make a
measurement of an object
without disturbing it. For
example, a thermometer
changes the temperature
of an object.
30.)What is the Heisenberg
uncertainty principle?
It
is impossible to know
precisely both the velocity
and position of a particle
at the same time.
The Schrodinger Wave Equation
32.) The Austrian physicist Edwin
Schrodinger derived an equation
that treated the electron as a
wave. What was significant about
his equation?
It worked equally well for
atoms of other elements
unlike Bohr’s model which
worked only for hydrogen.
33.) What is the name of the model
proposed in which electrons are
treated as waves?
The
quantum mechanical
model of the atom.
34.)What is the quantum mechanical
model of the atom?
It
limits the electrons energy
to certain values but makes
no attempt to describe the
path of the electron around
the nucleus
35.) What is an atomic orbital?
In
the quantum mechanical
model of the atom, the
orbital is the region in which
we expect to find an electron
90% of the time.
Orbitals are described by
electron clouds.
Hydrogen’s Atomic Orbital
36.)What are principal quantum
numbers?
 Principal
quantum numbers
represent the relative sizes and
energies of the atomic orbitals.
 As n increases, the size and
energy level increases.
37.) What are sublevels?
Each
principle energy level
has sublevels. The number of
sublevels is equal to the n
level.
 n = 1 has one sublevel
 n = 2 has two sublevels
 etc.
38.) What type of sublevels are
there?
The
sublevels are named s,
p, d, and f
Sublevel
s
p
Shape
spherical
energy level
lowest
dumbell
d
cloverleaf
f
complex
highest
39.)How many electrons does an
orbital contain?
two
40.) How many orbitals does each
sublevel contains?
Sublevel
orbitals
# e-
s
1
2
p
d
f
3
5
7
6
10
14
s
p
d
f
n=4
s
p
d
n= 3
s
p
n= 2
s
n=1
Principal Sublevel
Quantum types
#
#
orbitals
Total #
of
orbitals
Total #
electrons
1
s
1
1
2
2
s
p
1
3
4
8
3
s
P
d
s
p
d
f
1
3
5
1
3
5
7
9
18
16
32
4-7
Section 3 Electron Configurations
The arrangement of electrons in
atoms follows a few very specific
rules.
Ground State Electron Configurations
41.)What is an electron
configuration?
The
arrangement of electrons
in an atom
42.) What is ground state electron
configuration?
The most stable, lowest energy
arrangement of electrons.
43.) What are the three rules that
govern ground state electron
configurations.
1.) Aufbau principle
2.) Hund’s Rule
3.) Pauli Exclusion Principle
44.)What is the Aufbau principle?
Electrons
occupy the
lowest energy orbital
available.
German for building up
45.) In general what is the order of
increasing energy among energy
levels and sublevels?
In
general n=1 is lower than
n=2 . The order of increasing
order of sublevels is s, p, d, f.
Use the diagram
46.) What is the Pauli Exclusion
Principle?
A
maximum of two electrons
can occupy an orbital and they
must have opposite spins.
 The way to indicate two
electrons with opposite spin is
↓↑
47.) What is the Hund’s Rule?
Single electrons with the
same spin must occupy
equal energy orbitals
before additional electrons
with opposite spin can
occupy the same orbital.
48.)What are the two ways that
electron configurations can be
described?
 Orbital
diagrams – using boxes
 Mg
↓↑
1s
↓↑
2s
↓↑ ↓↑ ↓↑
2p
↓↑
3s
49.) What is noble gas configuration?
 Electron
configurations that are used
to show just the valence electrons.
The full inner core orbitals are
represented by the noble gas symbol
with the lower atomic number and
the electrons in the valence shell.
 K 1s22s22p63s23p64s1
 K [Ar] 4s1
50.) Exceptions to predicted
configurations.
 Cr
[Ar]4s13d5
 Cu [Ar]4s13d10
 It is more stable to “borrow” an
electron from the s orbital and put it
in the d orbital
Pg 160 practice problems
51.) Which electrons determine the
chemical properties of an element?
 The
valence electrons/ the outermost
electrons.
52. How does the number of valence
electrons compare to the family the
element is in?
Group 1 – 1 valence electron
 Group 2 – 2 valence electrons
 Group 13 – 3 valence electrons
 Group 14 – 4 valence electrons
 Group 15 – 5 valence electrons
 Group 16 - 6 valence electrons
 Group 17 – 7 valence electrons
 Group 18 – 8 valence electrons except for
hydrogen which only has two.
