Electrons in
Atoms
Basic Properties of Waves
Amplitude

Amplitude is the maximum distance
the particles of the medium carrying
the wave move away from their rest
positions.
Amplitude

Greater amplitude corresponds with
greater energy.
Wavelength

Wavelength (λ) is the distance
between two corresponding parts of a
wave.
Frequency

Frequency equals the number of
complete waves that pass a given
point in a certain amount of time.
Frequency and Wavelength

Frequency and wavelength are
inversely related, which means that as
one goes up the other goes down.
Frequency and Wavelength

Different wavelengths (and
frequencies) of light correspond
to different colors of light.
The Quantum Concept

In 1900, the German physicist Max
Planck began searching for an
explanation as he studied the light
emitted from heated objects.
The Quantum Concept
Matter can gain or lose energy
only in small, specific amounts
called quanta.
 That is, a quantum is the
minimum amount of energy that
can be gained or lost by an atom.

The Quantum Concept
While a beam of light has many
wavelike characteristics, it also
can be thought of as a stream of
tiny particles, or bundles of energy,
called photons.
 Thus, a photon is a particle of
electromagnetic radiation with no
mass that carries a quantum of
energy.

Energy and Frequency
Planck showed that the energy of
radiation increases as the
radiation’s frequency increases.
 Energy and frequency are directly
proportional.

Wave-Particle Duality of
Electrons
STOP HERE
The Bohr Model of the Atom

Niels Bohr produced a model of the
hydrogen atom based on
experimental observations.
The Bohr Model of the Atom

This model indicated that
1. An electron
circles the
nucleus only
in fixed
energy ranges
called orbits.
The Bohr Model of the Atom
2. An electron can neither gain nor lose
energy inside this orbit, but could
move up or down to another orbit.
The Bohr Model of the Atom
3. The lowest energy orbit is closest
to the nucleus.
Modern View
The atom has two
regions and is 3dimensional.
 The nucleus is at
the center and
contains the
protons and
neutrons.

Modern View

The electron cloud
is the region where
you might find an
electron (the region
of high probability)
and most of the
volume of an atom.
Wave-Particle Duality of
Electrons

Recall that electrons have wavelike properties, although they are
normally thought of as particles.
Particle Properties of
Electrons
There is always an integer number
of electrons orbiting the nucleus.
 Electrons jump
between orbitals
(orbits) in a
particle-like fashion.

Wave Properties of Electrons

The electrons do
not orbit the
nucleus in the
sense of a planet
orbiting the sun,
but instead exist as
standing waves.
Wave Properties of Electrons

The electrons are
never in a single point
location, although
there is an area of high
probability (the
electron cloud) of
interacting with the
electron at a single
point.
The Quantum
Mechanical Model

The lowest allowable energy state
of an atom is called its ground
state.
The Quantum
Mechanical Model
When an electron gains an amount
of energy equivalent to the energy
difference, it is said to be in an
excited state.
 The electron moves from its
ground state to a higher energy
level.

The Quantum
Mechanical Model
The Quantum
Mechanical Model

When the electron moves to a lower
energy level, it releases an amount
of energy equal to the energy
difference in these levels as
electromagnetic radiation (emissions
spectrum).
The Quantum
Mechanical Model

This electromagnetic radiation is
given off as photons.

The spectrum of light released from
excited atoms of an element is called
the emission spectrum of that
element.
• These are called
line spectra.
• Each is unique
to an element.
Ultraviolet
Visible
Infrared
The further electrons fall, the greater
the energy.
 This results in a higher frequency
because energy and frequency are
directly proportional.

Question
1. Use the Chemistry Reference
Tables to answer the following:
a) An electron falls from energy
level 5 to energy level 3.
What is the wavelength of the
light emitted?
(1282 nm)
Question
b) An electron falls from energy
level 6 to energy level 2.
What is the wavelength of the
light emitted?
(410 nm)
Question
c) An electron falls from energy
level 3 to energy level 1.
What type of electromagnetic
radiation is emitted (infrared,
visible or ultraviolet)?
(UV)
Question
d) An electron falls from energy
level 4 to energy level 2.
What type of electromagnetic
radiation is emitted (infrared,
visible or ultraviolet)?
(visible)
Question
e) An electron falls from energy
level 5 to energy level 2.
What color of visible light is
emitted?
(blue)
Question
f) An electron falls from energy
level 3 to energy level 2.
What color of visible light is
emitted?
(red)
The Quantum
Mechanical Model

The space around the nucleus of
an atom where the atom’s
electrons are found is called the
electron cloud.

A three-dimensional region
around the nucleus called an
atomic orbital describes the
electron’s probable location.
Energy Levels

In general, electrons reside in
principal energy levels.
Energy Levels

As the energy level number
increases, the orbital becomes
larger, the electron spends more
time farther from the nucleus, and
the atom’s energy level increases.
Sublevels
Principal energy levels contain
energy sublevels.
 Principal energy level 1 consists of
a single sublevel, principal energy
level 2 consists of two sublevels,
principal energy level 3 consists of
three sublevels, and so on.

Sublevels
Sublevels are labeled s, p, d, or f.
 The s sublevel can hold 2 electrons,
the p sublevel can hold 6 electrons,
the d sublevel can hold 10
electrons, and the f sublevel can
hold 14 electrons.

s block
p block
d block
f block
Orbitals
Sublevels contain orbitals.
 Each orbital may contain at most
two electrons.

s orbitals
One s orbital
for every
energy level
 The s orbital
is spherical
shaped.
 Called the 1s, 2s, 3s, etc… orbitals

p orbitals
Start at the second energy level
 p orbitals reside along 3 different
directions
 p orbitals have 3 different dumbbell
shapes

d orbitals
d orbitals start at the third energy level
 d orbitals have 5 different shapes

f orbitals
f orbitals start at the fourth energy
level
 f orbitals have seven different shapes

f orbitals
Summary
# of
shapes
(orbitals)
Max # of
electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
s block
p block
d block
f block
s1
s- block
s2
s2
Really have to include
helium.
 Helium has the properties
of the noble gases.

Transition Metals - d block
d1
d2
d3
d4
d5
d6
d7
d8
d9 d10
The p-block
p1 p2
p3
p4
p5
p6
f - block

inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

d orbitals fill up after previous
energy level so first d is 3d even
though it’s on row 4.
1
2
3
4
5
6
7
3d
1
2
3
4
5
6
7
4f
5f

f orbitals start filling at 4f
1
2
3
4
5
6
7

Each row (or period) is the energy
level for s and p orbitals.
Electron Configurations
Electron configurations
represent the way electrons are
arranged in atoms.
 Electrons enter the lowest energy
level first.

Electron Configurations
2. Determine the electron
configuration for phosphorus (P).

The atomic number of
phosphorus is 15, so we need to
account for 15 electrons.
Electron Configurations
• 1s2 2s2 2p6 3s2 3p3
• 12
2 electrons
4
10
15
electrons
• The 3p sublevel can hold 6
electrons, but only 3 are needed to
get the 15 electrons.
Electron Configurations
3. Determine the electron
configuration for chromium (Cr).

The atomic number of chromium
is 24, so we need to account for
24 electrons.
Electron Configurations
• 1s2 2s2 2p6 3s2 3p6 4s2 3d4
• 24
2 electrons
4
10
18
12
20
electrons
• The 3d orbital does not need all 10
electrons to get the 24 electrons.
Question
4. Write the electron configuration
for aluminum (Al).
(1s2 2s2 2p6 3s2 3p1)
Question
5. Write the electron configuration
for neon (Ne).
(1s2 2s2 2p6)
Question
6. Write the electron configuration
for calcium (Ca).
(1s2 2s2 2p6 3s2 3p6 4s2)
Question
7. Write the electron configuration
for iron (Fe).
(1s2 2s2 2p6 3s2 3p6 4s2 3d6)
Question
8. Write the electron configuration
for bromine (Br).
(1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5)
Electron Configuration and the
Bohr Model
Electron Configuration and the
Bohr Model
• Electron configuration was first
conceived of under the Bohr model
of the atom.
Electron Configuration and the
Bohr Model
Consider the
electron
configuration and
Bohr model image
for silicon (Si).
1s2 2s2 2p6 3s2 3p2
Electron Configuration and the
Bohr Model
1s2 2s2 2p6 3s2 3p2
There are 2
electrons in the
first energy level
(orbit).
Electron Configuration and the
Bohr Model
1s2 2s2 2p6 3s2 3p2
There are 8
electrons in the
second energy
level (orbit).
Electron Configuration and the
Bohr Model
1s2 2s2 2p6 3s2 3p2
There are 4
electrons in the
third energy level
(orbit).
STOP HERE
Orbital Diagrams
Remember sublevels contain
orbitals, and each orbital may
contain at most two electrons.
 The 2 electrons in an orbital must
have opposite “spins.”

Orbital Diagrams

In addition, when electrons occupy
orbitals of equal energy, they don’t
pair up with an electron of opposite
spin until they have to.
Orbital Diagrams
Ne
Example
9. Draw the orbital diagram for
oxygen (O).
The electron configuration for
oxygen is: 1s2 2s2 2p4
The orbital diagram is:
1s
2s
2p
Problem
10. Draw the orbital diagram for
silicon (Si).
The electron configuration for silicon
is: 1s2 2s2 2p6 3s2 3p2
The orbital diagram is:
1s
2s
2p
3s
3p
Problem
11. Draw the orbital diagram for potassium (K).
The electron configuration for is:
1s2 2s2 2p6 3s2 3p6 4s1
The orbital diagram is:
1s 2s
2p
3s
3p
4s
Using Electron Configurations
for Element Identification

To identify an element with a
given electron configuration, add
the superscript numbers
together and find the element with
that atomic number.
Example
12. Identify the element with the
following electron configuration:
1s2 2s2 2p6 3s1
Example

1s2 2s2 2p6 3s1

2 + 2 + 6 + 1 = 11

Element 11 is sodium (Na).
Question
13. Identify the element with the
following electron configuration:
1s2 2s2 2p6 3s2 3p4
(sulfur - S)
Question
14. Identify the element with the
following electron configuration:
1s2 2s2 2p6 3s2 3p6 4s2 3d9
(copper - Cu)
Question
15. Identify the element with the
following electron configuration:
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
(germanium - Ge)
Electron Configuration Using a
Noble Gas Abbreviation

In order to write this type of
configuration, find the noble gas
(from Group 8A) that comes
before the element in question.
Electron Configuration Using a
Noble Gas Abbreviation

Put the symbol for the noble gas
in brackets and then write the part
of the configuration that follows to
reach the desired element.
Example
16. Write the electron configuration
using a noble gas abbreviation
for magnesium (Mg).

Neon is the noble gas that
proceeds magnesium.
Example

Put neon’s symbol in brackets.

[Ne]

Now use the periodic table to
determine the rest of the
configuration.
1
s block
p block
2
3
d block
4
5
6
7
f block
Neonnoble
The
is in light
gas configuration
electron
blue; magnesium
configuration
additional
is 3s2is.
in bright
for
magnesium
yellow. is: [Ne] 3s2
Example
17. Write the electron configuration
using a noble gas abbreviation
for nickel (Ni).

Argon is the noble gas that
proceeds nickel.
Example

Put argon’s symbol in brackets.

[Ar]

Now use the periodic table to
determine the rest of the
configuration.
1
s block
p block
2
3
d block
4
5
6
7
f block
Argon
The
additional
noble
is ingas
lightconfiguration
electron
blue; nickel
configuration
is
is in
4s2 3d8.
bright
(Remember
for
nickel
yellow.
is: you
[Ar] subtract
4s2 3d8 1 from the d
sublevel row number.)
Question
18. Write the electron configuration
using a noble gas abbreviation
for fluorine (F).
([He] 2s2 2p5)
Question
19. Write the electron configuration
using a noble gas abbreviation
for silicon (Si).
([Ne] 3s2 3p2)
Question
20. Write the electron configuration
using a noble gas abbreviation
for zirconium (Zr).
([Kr] 5s2 4d2)
Valence Electrons
• The electrons in the outermost energy
level are called valence electrons.
• You can also use the periodic table as
a tool to predict the number of valence
electrons in any atom in Groups 1, 2,
13, 14, 15, 16, 17, and 18.
Valence Electrons
• All atoms in Group 1, like
hydrogen, have one valence
electron. Likewise, atoms in
Group 2 have two valence
electrons.
Valence Electrons
• All atoms in Group 13 have three
valence electrons.
• All atoms in Group 14 have four
valence electrons.
• All atoms in Group 15 have five
valence electrons.
Valence Electrons
• All atoms in Group 16 have six
valence electrons.
• All atoms in Group 17 have seven
valence electrons.
• All atoms in Group 18 have eight
valence electrons, except helium
which only has two.
Valence Electrons
• All atoms in Groups 3 – 11
(sublevels d and f ) have 2
valence electrons.
Question
21. How many valence electrons
does each of the following
elements have?
a) carbon (C)
(4)
b) bromine (Br) (7)
Question
c) iron (Fe)
(2)
d) potassium (K)
e) aluminum (Al)
(1)
(3)
Question
22. How many valence electrons
does each have based on the
following electron configuration?
a) [Kr] 5s2 4d2
(2)
b) [He] 2s2 2p5
(7)
Question
c) 1s2 2s2 2p6 3s2 3p6 4s2 3d9
d) 1s2 2s2 2p6 3s2 3p1
(2)
(3)
e) 1s2 2s2 2p6 3s2 3p6 4s2
(2)
Lewis Dot Diagrams
• Because valence electrons are so
important to the behavior of an
atom, it is useful to represent them
with symbols.
Lewis Dot Diagrams
• A Lewis dot diagram illustrates
valence electrons as dots (or other
small symbols) around the chemical
symbol of an element.
• Each dot represents one valence
electron.
• In the dot diagram, the element’s
symbol represents the core of
the atom—the nucleus plus all
the inner electrons.
Electron (Lewis) Dot
Diagrams

Write the symbol.

Put one dot for each
valence electron.

Don’t pair electrons until
you have to.
X
Question
23. Write a Lewis dot diagram for
chlorine.
Question
24. Write a Lewis dot diagram for
calcium.
Question
25. Write a Lewis dot diagram for
potassium.