Download Chapter 13 Electrons in Atoms

Chapter 13
Electrons in Atoms
Section 13.1
Models of the Atom
 OBJECTIVES:
Summarize the development of
atomic theory.
Explain the Quantum
Mechanical model and the theory
that electrons form an electron
“cloud”.
Greek Idea
Democritus and
Leucippus
 Matter is made up
of solid indivisible
particles
 John Dalton - one
type of atom for
each element

J. J. Thomson’s Model
Discovered electrons
 Atoms were made of
positive stuff
 Negative electron
floating around
 “Plum-Pudding”
model

Ernest Rutherford’s Model
Discovered dense
positive piece at
the center of the
atom- nucleus
 Electrons would
surround it
 Mostly empty
space
 “Nuclear model”

Niels Bohr’s Model







He had a question: Why don’t the electrons
fall into the nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Have different energies and therefore orbit at
different levels.
Cannot exist between orbits.
A quanta is the energy needed to jump to a
higher energy level- “Quantum Leap”
“Planetary model”
Bohr’s Planetary Model
The Quantum Mechanical
Model




Erwin Schrodinger derived an
equation that described the
energy and position of the
electrons in an atom
Things that are very small
behave differently from things big
enough to see.
The quantum mechanical model
is a mathematical solution
It is not like anything you can
see.
The Quantum Mechanical
Model
Has energy levels for
electrons.
 Orbits are not circular.
 They are not even ovals, they are
random three-dimensional shapes.
 It can only tell us the probability of
finding an electron a certain distance
from the nucleus.

The Quantum Mechanical
Model



The atom is found inside
a blurry “electron cloud”
Think of fan blades
spinning fast.
Electrons are moving so
fast that they create this
kind of blur. Except that
they are moving in a 3D
space.
Atomic Orbitals




Principal Quantum Number (n) = the energy
level of the electron.
Within each energy level, the complex math
of Schrodinger’s equation describes several
shapes.
These are called atomic orbitals - regions
where there is a high probability of finding an
electron.
Sublevels- like theater seats arranged in
sections
Summary
# of
Max
shapes electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
By Energy Level


First Energy Level

Third energy level
only s orbital
s, p, and d orbitals
only 2 electrons
18 total electrons
Second Energy
Level
s and p orbitals are
available
8 total electrons

Fourth energy level
s,p,d, and f orbitals
32 total electrons
By Energy Level
Any more than the fourth and not all the
orbitals will fill up.
 You simply run out of electrons
 The orbitals do not fill up in a neat
order.
 The energy levels overlap
 Lowest energy fill first.

Section 13.2
Electron Arrangement in Atoms
 OBJECTIVES:
Apply the aufbau principle, the
Pauli exclusion principle, and
Hund’s rule in writing the electron
configurations of elements.
Section 13.2
Electron Arrangement in Atoms
 OBJECTIVES:
Explain why the electron
configurations for some elements
differ from those assigned using
the aufbau principle.
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2p
2s
1s
Aufbau diagram - page 367
Electron Configurations
The way electrons are arranged in
atoms.
 Aufbau principle- electrons enter the
lowest energy first.
 This causes difficulties because of the
overlap of orbitals of different energies.
 Pauli Exclusion Principle- at most 2
electrons per orbital - different spins

Electron Configuration
 Hund’s
Rule- When electrons
occupy orbitals of equal energy
they don’t pair up until they have to.
 Let’s determine the electron
configuration for Phosphorus
 Need to account for 15 electrons
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3s
2s
1s
4f
3d
4s
3p
5f
The first two electrons
go into the 1s orbital
2p
 Notice the opposite
spins
 only 13 more to go...

Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
5f
4f
3d
4s
3p
3s
2p
2s
1s
The next electrons
go into the 2s orbital
 only 11 more...

Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2p
2s
1s
• The next electrons go
into the 2p orbital
• only 5 more...
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2p
2s
1s
• The next electrons go
into the 3s orbital
• only 3 more...
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into
separate shapes
3 unpaired electrons
= 1s22s22p63s23p3
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
Now do Oxygen.
5f
4f
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
1s2
6d
5d
4d
3d
5f
4f
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
1s22s2
6d
5d
4d
3d
5f
4f
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
1s22s22p4
6d
5d
4d
3d
5f
4f
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
6d
5d
4d
5f
4f
3d
Now Bromine
2p
2s
1s22s22p63s23p64s23d104p5
1s
Exceptional Electron
Configurations
Orbitals fill in order
 Lowest
energy to higher energy.
 Adding electrons can change the
energy of the orbital.
 Half filled orbitals have a lower
energy.
 Makes them more stable.
 Changes the filling order
Write these electron
configurations
Titanium - 22 electrons
1s22s22p63s23p64s23d2
 Vanadium - 23 electrons
1s22s22p63s23p64s23d3
 Chromium - 24 electrons
1s22s22p63s23p64s23d4 expected
But this is wrong!!

Chromium is actually:
 1s22s22p63s23p64s13d5
 Why?
 This
gives us two half filled orbitals.
 Slightly lower in energy.
 The same principal applies to
copper.
Copper’s electron
configuration
Copper has 29 electrons so we expect:
1s22s22p63s23p64s23d9
 But the actual configuration is:
 1s22s22p63s23p64s13d10
 This gives one filled orbital and one half
filled orbital.
 Remember these exceptions: d4, d9

Section 13.3
Physics and the Quantum
Mechanical Model

OBJECTIVES:
Calculate the wavelength,
frequency, or energy of light, given
two of these values.
Section 13.3
Physics and the Quantum
Mechanical Model

OBJECTIVES:
Explain the origin of the atomic
emission spectrum of an element.
If the light is not white
By heating a gas
with electricity we
can get it to give
off colors.
 Passing this light
through a prism
does something
different.

Atomic Spectrum
Each element
gives off its own
characteristic
colors.
 Can be used to
identify the atom.
 How we know
what stars are
made of.

• These are called
discontinuous
spectra, or line
spectra
• unique to each
element.
• These are
emission spectra
• The light is emitted
given off
• Sample 13-2 p.375
Explanation of atomic spectra
 When
we write electron
configurations, we are writing the
lowest energy.
 The energy level, and where the
electron starts from, is called it’s
ground state- the lowest energy
level.
Changing the energy
 Let’s
look at a hydrogen atom
Changing the energy

Heat or electricity or light can move the
electron up energy levels (“excited”)
Changing the energy

As the electron falls back to ground
state, it gives the energy back as light
Changing the energy
May fall down in steps
 Each with a different energy

Ultraviolet
Visible
Infrared
 Further they fall, more energy, higher
frequency.
 This is simplified
 the orbitals also have different energies
inside energy levels
 All the electrons can move around.
The physics of the very small
 Quantum
mechanics explains how
the very small behaves.
 Classic physics is what you get
when you add up the effects of
millions of packages.
 Quantum mechanics is based on
probability
Heisenberg Uncertainty
Principle
 -It
is impossible to know exactly the
location and velocity of a particle.
 The better we know one, the less
we know the other.
 Measuring changes the properties.
 Instead, analyze interactions with
other particles
More obvious with the very
small
 To
measure where a electron is, we
use light.
 But the light moves the electron
 And hitting the electron changes the
frequency of the light.
Before
After
Photon
Moving
Electron
Photon
changes
wavelength
Electron
Changes
Velocity
Fig. 13.19, p. 382