# Chapter 13 Electrons in Atoms by HC111123011234

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```									   Chapter 13
Electrons in Atoms
Section 13.1
Models of the Atom
 OBJECTIVES:
Summarize the development of
atomic theory.
Section 13.1
Models of the Atom
 OBJECTIVES:
Explain the significance of
quantized energies of electrons
as they relate to the quantum
mechanical model of the atom.
Greek Idea
 Democritus
of solid indivisible
particles
 John Dalton - one
type of atom for
each element
J. J. Thomson’s Model
 Discovered electrons
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.
 Amounts of energy separate one
level from another.
 “Planetary model”
Bohr’s planetary model
 Energy level of an electron
 analogous to the rungs of a ladder
 electron cannot exist between energy
levels, just like you can’t stand
 Quantum of energy required to move
to the next highest level
The Quantum Mechanical
Model
 Energy is quantized. It comes in chunks.
 A quanta is the amount of energy needed to
move from one energy level to another.
 Since the energy of an atom is never “in
between” there must be a quantum leap in
energy.
 Erwin Schrodinger derived an equation that
described the energy and position of the
electrons in an atom
The Quantum Mechanical
Model
 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.
 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”
 A area where there is
a chance of finding
an electron.
 Draw a line at 90 %
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
Two
representations
of the hydrogen
1s, 2s, and 3s
orbitals.
Representation of the 2p orbitals. (a) The
electron probability distributed for a 2p
orbital. (b) The boundary surface
representations of all three 2p orbitals.
Representation of the 3d orbitals.
Representation of the 4f orbitals in
terms of their boundary surface.
Summary
# of   Max         Starts at
shapes electrons   energy level
s    1        2            1

p    3        6            2

d    5        10           3

f    7        14           4
By Energy Level
 First Energy Level    Second Energy
 only s orbital         Level
 only 2 electrons      s and p orbitals

 1s
2                   are available
 2 in s, 6 in p
2 6
 2s 2p
 8 total electrons
By Energy Level
 Third energy level     Fourth energy
 s, p, and d             level
orbitals               s,p,d, and f
 2 in s, 6 in p, and     orbitals
10 in d                2 in s, 6 in p, 10
2 6 10
 3s 3p 3d                in d, ahd 14 in f
 4s
24p64d104f14
 18 total electrons
 32 total electrons
By Energy Level
 Any more than            The orbitals do
the fourth and not        not fill up in a
all the orbitals will     neat order.
fill up.                 The energy levels
 You simply run            overlap
out of electrons         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.
7p         6d
7s   6p                      5f
5d
6s   5p                      4f
4d
5s
Increasing energy

4p
4s              3d
3p
3s
2p
2s

Aufbau diagram - page 367
1s
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
7p          6d
7s   6p                        5f
5d
6s   5p                        4f
4d
5s
Increasing energy

4p
4s               3d
3p   The first two electrons

3s
go into the 1s orbital
2p
2s       Notice the opposite
spins
 only 13 more to go...
1s
7p           6d
7s   6p                          5f
5d
6s   5p                          4f
4d
5s
Increasing energy

4p
4s                3d
3p
3s
2p    The next electrons
2s          go into the 2s orbital
 only 11 more...

1s
7p          6d
7s   6p                        5f
5d
6s   5p                        4f
4d
5s
Increasing energy

4p
4s               3d
3p
3s
2p   • The next electrons go
2s          into the 2p orbital
• only 5 more...
1s
7p          6d
7s   6p                         5f
5d
6s   5p                         4f
4d
5s
Increasing energy

4p
4s               3d
3p
3s
2p   • The next electrons go
2s          into the 3s orbital
• only 3 more...
1s
7p              6d
7s   6p                              5f
5d
6s   5p                              4f
4d
5s
Increasing energy

4p
4s                    3d
3p •     The last three electrons
3s            go into the 3p orbitals.
2p •     They each go into
2s            separate shapes
•   3 unpaired electrons
1s        •   = 1s22s22p63s23p3
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
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
Chemical Physics and the
Quantum Mechanical Model
   OBJECTIVES:
Calculate the wavelength,
frequency, or energy of light, given
two of these values.
Section 13.3
Chemical Physics and the
Quantum Mechanical Model
   OBJECTIVES:
Explain the origin of the atomic
emission spectrum of an element.
Light
 The  study of light led to the
development of the quantum
mechanical model.
 Light is a kind of electromagnetic
many kinds of waves
 All move at 3.00 x 108 m/s = c
Parts of a wave
Crest
Wavelength

Amplitude
Origin

Trough
Parts of Wave - p.372
   Origin - the base line of the energy.
   Crest - high point on a wave
   Trough - Low point on a wave
   Amplitude - distance from origin to trough (-)
or crest (+)
   Wavelength - distance from crest to crest
   Wavelength is abbreviated by the Greek letter
lambda = l
Frequency
 The number of waves that pass a
given point per second.
 Units: cycles/sec or hertz (Hz or sec-1)
 Abbreviated by Greek letter nu = n

c = ln
Frequency and wavelength
 Are inversely related
 As one goes up the other goes down.
 Different frequencies of light are
different colors of light.
 There are a wide variety of frequencies
 The whole range is called a spectrum,
Fig. 13.10, page 373
Low                               High
energy                            energy

Radio Micro Infrared    Ultra- X-     Gamma
waves waves .           violet Rays Rays
Low                                High
Frequency                          Frequency
Long                             Short
Wavelength                       Wavelength
Visible Light
Prism
 White light is
colors of the
visible spectrum.
 Passing it through
a prism separates
it.
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
• 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
Light is a Particle
 Energy   is quantized.
 Light is energy
 Light must be quantized
 These smallest pieces of light are
called photons.
 Photoelectric effect?
 Energy & frequency: directly related.
Energy and frequency
E  =hxn
 E is the energy of the photon
 n is the frequency
 h is Planck’s constant
 h = 6.6262 x 10 -34 Joules x sec.
 joule is the metric unit of Energy
The Math in Chapter 13

2  equations so
far:
 c = ln
 E = hn
 Know these!
Examples
 What  is the wavelength of blue light
with a frequency of 8.3 x 1015 hz?
 What is the frequency of red light
with a wavelength of 4.2 x 10 -5 m?

 What is the energy of a photon of
each of the above?
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.
What is light?
 Light is a particle - it comes in chunks.
 Light is a wave- we can measure its
wavelength and it behaves as a wave
2
 If we combine E=mc , c=ln, E = 1/2
mv2 and E = hn
 We can get: l = h/mv
 called de Broglie’s equation
 Calculates the wavelength of a particle.
Sample problem
 What  is the approximate mass of a
particle having a wavelength of 10-7
meters, and a speed of 1 m/s?
Use l = h/mv
= 6.6 x 10-27
(Note: 1 J = N x m; 1 N = 1 kg x
m/s2
Matter is a Wave
 Does  not seem to apply to large
objects
 Things bigger than an atom
 A baseball has a wavelength of
about 10-32 m when moving 30 m/s
 Anelectron at the same speed has
a wavelength of 10-3 cm
 Big   enough to measure.
The chemical 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
 Itis 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.
other particles
More obvious observations
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
Photon           changes
wavelength

Moving                  Electron
Electron                    Changes
velocity
Fig. 13.19, p. 382

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