Download Electrons in Atoms

Chapter 5
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Scientific Models
Models are things used to represent real
phenomena.
simplify and explain complex realities.
can take many forms
scale models, e.g. a globe
mathematical models, e.g. PV = k
computer models, etc., e.g. weather predictions
Rutherford’s Model
It explained much about
the structure
Nucleus: positive,
very dense, most of
atom’s mass
Electrons: outside the
nucleus
Empty space: most of
the volume of the
atom
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It could not explain chemical behavior of
elements, such as….
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Why did elements give off light when heated?
Why did one element react with another to form
a new compound?
It could not explain why electrons did not
fall into the nucleus
The Bohr Model
Neils Bohr
Danish Physicist
1913: Proposed
new model of the
atom
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Bohr proposed that an electron is found only in specific
circular paths, or orbits, around the nucleus.
Each possible electron orbit in Bohr’s model has a fixed
energy.
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The fixed energies an electron can have are called energy
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levels.
Electrons are restricted to the “allowed” energy
levels
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Higher energy levels are farther away from the nucleus
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A quantum of energy is the amount of energy required to
move an electron from one energy level to another energy level.
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Energy levels are like
rungs on a ladder
Higher energy levels
are closer together
Takes less energy to
change between higher
levels
The Bohr Model
When
electrons
absorb exactly the
right quanta of
energy….
They jump to
higher energy level
When it jumps
back down…
It gives off (emits)
the same energy as
light.
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Bohr’s planetary model
only worked for
hydrogen
But it could not explain
motion of electrons
Schrödinger and others
developed a new
mathematic model of
the atom….
Called the quantum
mechanical model
ERWIN SCHRÖDINGER
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Like Bohr’s model, electrons are restricted to
certain energy levels
Unlike Bohr’s model, the exact pathway of the
electron is uncertain
Locations of electrons are uncertain, and
described terms of probability….
i.e. the likelihood of finding the electron at a
given point in time
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Electrons are found within an “electron cloud”
outside the nucleus
The electron
cloud is more
dense where the
probability of
finding the
electron is high.
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A spinning fan blade
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Forms a ‘fuzzy’
image
You know the fan
blade is within the
fuzzy region, but at
any point in time you
don’t know exactly
where it is
Electrons are located
in regions of
probability called
“orbitals”
Quantum
Number
Defines
Describes
Values
1st Principal
Energy Level
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n = 1 to 7
2nd Angular
Momentum
Energy Sublevel
Shape
s, p, d, f
3rd Magnetic
Orbital
3-D orientation
x, y, z, etc.
4th Spin
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Magnetic spin
+1/2 or -1/2
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Orbitals are represented
by “electron density
maps”
Probability is
represented by the
density of color
The more probable
location of the electron
is in the darker blue
region
AN “s” ORBITAL
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Regions of space in which there is a high
probability of finding an electron
Various types of orbitals exist, depending upon
the sublevel
S sublevels have one orbital
P sublevels have 3 orbitals each
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d sublevels have 5 orbitals each
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f sublevels have 7 orbitals each
Energy
Level
# Sublevels
# Orbitals
Electron
capacity
n
n
n2
2n2
1
1
1
2
2
2
4
8
3
3
9
18
4
4
16
32
5
5
25
50
6
6
36
72
7
7
49
98
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Each orbital can contain up to 2 electrons!
Sublevel
# Orbitals per
sublevel
Electron capacity
per sublevel
s
1
2
p
3
6
d
5
10
f
7
14
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In most natural phenomena, change trends
toward lower energy
Systems are more stable when they have less
energy.
Electrons also tend to arrange themselves in
their lowest energy states.
The arrangement of electrons within an atom is
called an electron configuration.
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Three rules are used to determine electron
configurations
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
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Electrons occupy the lowest energy level first
This diagram is known as an electron orbital
diagram
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4th quantum number is the “spin” number
Electrons “spin”, either clockwise & counterclockwise
Spin is symbolized ↑ or ↓
PEP says….
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Two electrons in the same orbital must have
opposite spins. Therefore ….
No two electrons in an atom can have the same
identical set of 4 quantum numbers.
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Electrons fill orbitals within a sublevel such
that they have maximum number of unpaired
spins
This is because they have the lowest energy
this way
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Determine the number of electrons in the
diagram. How?
Begin filling orbitals at the lowest energy level
(Aufbau principle)
Continue filling, applying Hund’s rule
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All “up” spins
Follow by “down” spins
Stop when you have assigned all the electrons
to orbitals
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A shorthand way for writing orbital diagrams
Write the energy level, sublevel, and number of
electrons in the sublevel
Li
1s2 2s1
C
1s2 2s2 2p2
N
1s2 2s2 2p3
O
1s2 2s2 2p4
F
1s2 2s2 2p5
Ne
1s2 2s2 2p6
Na
1s2 2s2 2p6 3s1
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Periods (rows) in the PT correspond to energy level
Certain groups (columns) correspond to the sublevels
(s, p, d, f) (see page 166)
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Also called “noble gas notation”
An element’s electron configuration contains
the e-config of a noble gas (group VIIIA, 18)
Begin with the preceding noble gas
Then complete the e-config
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Transition elements (groups 3-12) tend to
prefer half-filled or completely filled d-orbitals
at the expense of the s-orbital.
For chromium, you would expect
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but in fact one of the 4s electrons is promoted
to 3d, resulting in
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….4s2 3d4,
….4s1 3d5
Try copper….
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Much of what is known about the atom is due
to the study of light
Light has properties of waves
Waves have amplitude, wavelength, and
frequency
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Inversely proportional
c  
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c = speed of light = 3.00 x 108 m/s
(constant)
lambda = wavelength (meters)
nu = frequency (Hertz, Hz, s-1)
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Visible light is a small portion of the electromagnetic spectrum
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All EM radiation travels at the same speed
c = 3.00 x 108 m/s
EM radiation varies in wavelength and
frequency
Longer wavelength → Lower frequency
Shorter wavelength → Higher frequency
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Light separates into different colors
(wavelengths) when it passes through a prism
It is a continuous spectrum
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Electrons of an element can absorb energy and
emit the energy as EM radiation
These emission spectra are not continuous
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Each element has a unique emission spectra
Like a bar code for an element
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Electron at ground
state absorbs a
quantum of energy
Excited electron
returns to ground
state, emitting the
quantum as light
Frequency of the light
is directly
proportional to the
energy change of the
electron
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Lyman Series is in
the UV range
Balmer series is
visible
Paschen series is
in IR range
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Einstein determined that light behaved like a
particle
“Particle” of light is the photon
Photon is a quantum of light
So light can behave as a wave and a particle,
which is it?
Both
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If light (a wave) can behave as a particle, can a
particle behave as a wave?
Yes
So electrons can be thought of as waves.
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Uncertainty principle
It is not possible to know the location and
momentum (speed) of an electron at the same
time
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Schrodinger equation
Mathematically described sublevels and
orbitals