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Chapter 5
Electrons in Atoms
Mr. Samaniego
Lawndale High School
Summary of Atomic Theory
Year
400BC
Event
Democritus proposes idea of atom
1808
Dalton develops Atomic Theory
1897
Thomson uses cathode ray to
discover electron
1916
Millikan measures the mass of an e-
1919
Rutherford uses gold foil experiment
to discover nucleus
Structure Of An Atom

So by this point, we know that protons and
neutrons are located in the nucleus and
electrons are around the outside of the nucleus
Section 5.1 – Models of the Atom
In 1897 J. J. Thomson discovered the electron
Observed that a magnet
deflected the straight paths of
the cathode rays
Atoms were known to be electrically neutral
which meant that there had to be some
positively charged matter to balance the
negative charges
Ernest Rutherford’s
experiment disproved the
plum pudding model of the
atom and suggested that
there was a positively
charged nucleus because
most of the alpha particles
went straight through the
gold foil
BUT, Rutherford’s
atomic model could
not explain the
chemical properties
of elements
The Bohr Model (Niels Bohr)
In 1913, Niels Bohr came up with a new
model (Bohr was a student of Rutherford)
He noticed that light given
out when atoms were
heated always had specific
amounts of energy, so
Niels Bohr proposed a
model that electrons in an
atom must be orbiting the
nucleus and can reside only
in fixed energy levels
Energy Levels
Each of the electrons in Bohr’s
model has a fixed amount of
energy called energy levels
This is similar to steps of a ladder
(can climb up the ladder, but cannot
step in between the steps)

Quantum
is the amount of energy
required to move an electron from
one energy level to another
The
further away from the nucleus, the more
energy the electron has
The Quantum Mechanics Model
(Erwin Schrodinger)

While Rutherford’s model
described the path the electron
moves, Erwin Schrodinger solved
mathematical equations to
describe the behavior of electron

Similar to Bohr’s model, Schrodinger
describes the energy of electrons with
certain values but does not involve an exact
path the electron takes around the nucleus
The Quantum Mechanics View
of the Atom (Schrodinger)
The quantum
mechanical model does
not describe the exact
path an electron takes
around the nucleus, but
determines the
probability of finding
an electron in a certain
area
Quantum Mechanical Model



In this model, electrons move similar to a
rotating propeller blade
You cannot tell its precise location at any instant
because it’s a blurry region, but you have
information regarding the probability of finding an
electron within a certain volume of space
Similar to a fuzzy cloud…the probability of finding an
electron is higher where the cloud is more dense
Atomic Orbitals

An Atomic Orbital is a region of space
where there is a high probability of
finding an electron
Each
energy sublevel
corresponds to an orbital
of different shape
describing where the
electron is likely to be
found (there are 4
different types of shapes)
Shapes of Orbitals
Shapes of Orbitals
Shapes of Orbitals
Chapter 5.2 – Electron
Arrangement in Atoms
Each electron in an atom is assigned a set of
four quantum numbers. These help to
determine the highest probability of finding the
electrons.
Three of these numbers (n, l, m) give the
location of the electron

The
fourth (s) describes the orientation of an electron
in an orbital.
Quantum letters can be thought of like the
numbers and letters on a concert ticket
Labeling Electrons in Atoms
Probable Location
of electron
Probability
Probable location of
Finding Beyonce or Taylor
Energy level (n)
High Probability
Hotel Name
Sublevel (l)
Higher
Probability
Highest
probability
Hotel Floor
Number
Hotel Room
Number
Orbital (m)
Electron Configuration
Sublevel
# of Orbitals
Available
# of Electrons
Available
s
p
d
1
3
5
2
6
10
f
7
14
Electron Configurations
Filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d,
4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p


Example He = 2 electrons
1s2
 Example Li = 3 electrons
1s22s1

Example B = 5 electrons
1s22s22p1
Practice Problems
Write electron configurations for the following atoms
1.
2.
3.
4.
Li
N
Be
C
5. P
6. Si
7. Mg
8. Al
Electron Configurations can be
written in terms of noble gases
To save space, configurations can be written
in terms of noble gases


Example 1: Ne = 1s22s22p6
S = 1s22s22p63s23p4
Or
S=
[Ne] 3s23p4
Example 2: Ar = 1s22s22p63s23p6
Mn = 1s22s22p63s23p64s23d5
Mn =
[Ar]
4s23d5
Reading the Periodic Table
Locating Electrons in Atoms
So far we have discussed 3 quantum numbers

n= principal quantum level (principal energy level)

l= Sublevel

m = magnetic quantum number (shape of orbitals)
1s2
n
Number of electrons in sublevel
l
s = spin
When an electron moves, it generates a
magnetic field.
 s describes the direction an electron spins


They must spin in opposite directions

Spin= up
down

There are two values of s: +1/2 and -1/2
Orbital Diagrams

The electron
configuration gives the
number of electrons in
each sublevel but does
not show how the
orbital of a sublevel are
occupied by the
electrons
Orbital Diagrams

Used to show how electrons are
distributed within sublevels
Each orbital is represented by a box and
each electron is represented by an arrow
 Notice that each box is drawn higher than
the last set because it has more energy
Example: Boron 1s22s22p1

1s
2s
2p
Orbital Diagrams
Steps to writing orbital diagrams:ex F (Z=9)
1.
Write the electron configuration
1s22s22p5
2. Construct an orbital filling diagram using boxes
for each orbital
1s 2s
2p
3. Use arrows to represent the electrons in each
orbital.
1s
2s
2p
Rule #1 - Aufbau Principle


Electrons must occupy the orbital with
the lowest energy first
Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Rule #2 - Pauli Exclusion Principle



Orbitals can only have two electrons max
The 2 electrons must have opposite spins
Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Rule #3 - Hund’s Rule


Orbitals of equal energy are each occupied
by one electron before any pairing occurs
Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Draw orbital diagrams for these
elements
1.
2.
3.
4.
Li
N
Be
C
5.
6.
7.
8.
P
Si
Mg
Al
Section 5.3 - Atomic Spectra


When atoms absorb energy, electrons
move into higher energy levels (excited
state)
When these electrons return to their lower
energy levels, they lose energy by emitting
light
 Atomic Emission Spectrum – the discrete
lines representing the frequencies of
light emitted by an element
Calculating Wavelength of Light
c=
c = speed of light (3 x 108 m/s2)
 = wavelength (called lambda)
 = frequency (called nu)
Practice
1. Calculate the wavelength of a yellow light
if the frequency is 5.10 x 1014 sec-1 or Hz.
Answer = 5.88 x 10-7m
2. What is the wavelength of 1.50 x 1013 sec-1?
Answer = 2.00 x 10-5m
3. What frequency is radiation with a wavelength
of 5.00 x 10-8m?
Answer = 6.00 x 1015 sec-1 or Hertz
Atomic Spectra



Each discrete line in an emission
spectrum corresponds to one exact
frequency of light emitted by the atom
Ground State – lowest possible energy of
the electron in the Bohr model
The light emitted by an electron moving
from higher to a lower energy level has
a frequency directly proportional to the
energy change of the electron
Homework
Chapter 5 Assessment Page 148
#’s 22-24, 27, 29, 30-39,
50-53, 57, 60, 68, 70-72