Transcript Chapter 5 ppt

Chapter 5
Section 5.1
Physics and the Quantum
Mechanical Model

OBJECTIVES:
•
Describe the relationship
between the wavelength and
frequency of light.
Section 5.1
Physics and the Quantum
Mechanical Model

OBJECTIVES:
•
Identify the source of atomic
emission spectra.
Section 5.1
Physics and the Quantum
Mechanical Model

OBJECTIVES:
•
Explain how the frequencies of
emitted light are related to
changes in electron energies.
Section 5.1
Physics and the Quantum
Mechanical Model

OBJECTIVES:
•
Distinguish between quantum
mechanics and classical
mechanics.
Light





The study of light led to the development of
the quantum mechanical model.
Light is a kind of electromagnetic radiation.
Electromagnetic radiation includes many
types: gamma rays, x-rays, radio waves…
Speed of light = 2.998 x 108 m/s, and is
abbreviated “c”
All electromagnetic radiation travels at this
same rate when measured in a vacuum
- Page 139
“R O Y
Frequency Increases
Wavelength Longer
G
B I
V”
Parts of a wave
Crest
Wavelength
Amplitude
Origin
Trough
Electromagnetic radiation propagates through
space as a wave moving at the speed of light.
Equation:
c =
c = speed of light, a constant (2.998 x 108 m/s)
 (lambda) = wavelength, in meters
 (nu) = frequency, in units of hertz (hz or sec-1)
Wavelength and Frequency

Are inversely related
•
As one goes up the other goes down.
Different frequencies of light are
different colors of light.
 There is a wide variety of frequencies
 The whole range is called a spectrum

Low
Energy
High
Energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
Low
High
Frequency
Frequency
Long
Short
Wavelength
Visible Light Wavelength
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High Frequency
=
High ENERGY
Wavelength Table

Water and sound waves transfer energy
from one place to another- they require
a medium through which to travel. They
are mechanical waves.
.
Nature of Electromagnetic Waves




They are Transverse waves without a medium.
(They can travel through empty space)
They travel as vibrations in electrical and
magnetic fields.
Have some magnetic and some electrical
properties to them.
Speed of electromagnetic waves = 300,000,000
meters/second (Takes light 8 minutes to move
from the sun to earth {150 million miles} at this
speed.)
When an electric field changes, so does the
magnetic field. The changing magnetic field
causes the electric field to change. When one
field vibrates—so does the other.
 RESULT-An electromagnetic wave.

Waves or Particles

Electromagnetic radiation has properties of waves
but also can be thought of as a stream of
particles. Example: Light

Light as a wave: Light behaves as a transverse
wave which we can filter using polarized lenses.

Light as particles (photons)

When directed at a substance light can knock
electrons off of a substance (Photoelectric effect)
Measuring the electromagnetic
spectrum


You actually know more about it than you may
think! The electromagnetic (EM) spectrum is just
a name that scientists give a bunch of types of
radiation when they want to talk about them as a
group.
Radiation is energy that travels and spreads out
as it goes-- visible light that comes from a lamp
in your house or radio waves that come from a
radio station are two types of electromagnetic
radiation.
Measuring the electromagnetic
spectrum
Other examples of EM radiation are
microwaves, infrared and ultraviolet light,
X-rays and gamma-rays.
 Hotter, more energetic objects and events
create higher energy radiation than cool
objects.
 Only extremely hot objects or particles
moving at very high velocities can create
high-energy radiation like X-rays and
gamma-rays.

B. Waves of the Electromagnetic Spectrum

Electromagnetic Spectrum—name for the range of
electromagnetic waves when placed in order of increasing
frequency

Click here (Animation—Size of EMwaves)
RADIO
WAVES
INFRARED
RAYS
MICROWAVES
ULTRAVIOLET
RAYS
VISIBLE LIGHT
GAMMA
RAYS
X-RAYS
Here are the different types of
radiation in the EM spectrum,
in order from lowest energy to
highest:
RADIO WAVES
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A. Have the longest wavelengths and lowest
frequencies of all the electromagnetic waves.
B. A radio picks up radio waves through an
antenna and converts it to sound waves.
(Garage doors and wireless networks)
C. Each radio station in an area broadcasts at a
different frequency. # on radio dial tells frequency.
D. MRI (MAGNETIC RESONACE IMAGING)

Uses Short wave radio waves with a magnet to create
an image
MRI of the Brain
Radio

Radio: yes, this is the same kind of energy
that radio stations emit into the air for your
boom box to capture and turn into your
favorite tunes. But radio waves are also
emitted by other things ... such as stars
and gases in space. You may not be able
to dance to what these objects emit, but
you can use it to learn what they are made
of.
AM=Amplitude modulation—waves bounce off ionosphere can
pick up stations from different cities.
(535kHz-1605kHz= vibrate at 535 to 1605 thousand times/second)
+
FM=Frequency modulation—waves travel in a straight line &
through the ionosphere--lose reception when you travel out of range.
(88MHz-108MHz = vibrate at 88million to 108million times/second)
+
Bands of Radio/TV/Microwaves
MICROWAVES

Microwaves—have the shortest wavelengths
and the highest frequency of the radio waves.

Used in microwave ovens.
 Waves
transfer energy to the water in the food causing
them to vibrate which in turn transfers energy in the
form of heat to the food.
Used by cell phones and pagers.
 RADAR (Radio Detection and Ranging)

 Used
to find the speed of an object by sending out radio
waves and measuring the time it takes them to return.
MICROWAVES

Microwaves: they will cook your popcorn in
just a few minutes! In space, microwaves
are used by astronomers to learn about
the structure of nearby galaxies, including
our own Milky Way!
INFRARED RAYS
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Infrared= below red
Shorter wavelength and higher frequency than
microwaves.
You can feel the longest ones as warmth on your
skin
Heat lamps give off infrared waves.
Warm objects give off more heat energy than cool
objects.
Thermogram—a picture that shows regions of
different temperatures in the body. Temperatures
are calculated by the amount of infrared radiation
given off. Therefore people give off infrared rays.
INFRARED RAYS

Infrared: we often think of this as being the
same thing as 'heat', because it makes our
skin feel warm. In space, IR light maps the
dust between stars.
VISIBLE LIGHT
Shorter wavelength and higher frequency than
infrared rays.
 Electromagnetic waves we can see.
 Longest wavelength= red light
 Shortest wavelength= violet (purple) light
 When light enters a new medium it bends
(refracts). Each wavelength bends a different
amount allowing white light to separate into
it’s various colors ROYGBIV.

VISIBLE LIGHT

Visible: yes, this is the part that our eyes
see. Visible radiation is emitted by
everything from fireflies to light bulbs to
stars ... also by fast-moving particles
hitting other particles.
ULTRAVIOLET RAYS
Shorter wavelength and higher frequency than
visible light
 Carry more energy than visible light
 Used to kill bacteria. (Sterilization of
equipment)
 Causes your skin to produce vitamin D (good
for teeth and bones)
 Used to treat jaundice ( in some new born
babies.
 Too much can cause skin cancer.
 Use sun block to protect against (UV rays)

ULTRAVIOLET RAYS

Ultraviolet: we know that the Sun is a
source of ultraviolet (or UV) radiation,
because it is the UV rays that cause our
skin to burn! Stars and other "hot" objects
in space emit UV radiation.
X- RAYS

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Shorter wavelength and higher frequency than UVrays
Carry a great amount of energy
Can penetrate most matter.
Bones and teeth absorb x-rays. (The light part of an
x-ray image indicates a place where the x-ray was
absorbed)
Too much exposure can cause cancer


(lead vest at dentist protects organs from unnecessary
exposure)
Used by engineers to check for tiny cracks in
structures.

The rays pass through the cracks and the cracks appear
dark on film.
Type:
JPG
GAMMA RAYS
Shorter wavelength and higher frequency than
X-rays
 Carry the greatest amount of energy and
penetrate the most.
 Used in radiation treatment to kill cancer
cells.
 Can be very harmful if not used correctly.
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GAMMA RAYS
•
•
•
Gamma-rays: radioactive materials (some
natural and others made by man in things like
nuclear power plants) can emit gamma-rays.
Big particle accelerators that scientists use to
help them understand what matter is made of
can sometimes generate gamma-rays.
But the biggest gamma-ray generator of all is
the Universe! It makes gamma radiation in all
kinds of ways.
Brief Summary
A. All electromagnetic waves travel at the
same speed. (300,000,000 meters/second
in a vacuum.
 B. They all have different wavelength and
different frequencies.

Long wavelength-lowest frequency
 Short wavelength highest frequency
 The higher the frequency the higher the
energy.

Section 5.2
Models of the Atom
 OBJECTIVES:
•
Identify the inadequacies in the
Rutherford atomic model.
Section 5.2
Models of the Atom
 OBJECTIVES:
•
Identify the new proposal in the
Bohr model of the atom.
Section 5.2
Models of the Atom

OBJECTIVES:
•
Describe the energies and
positions of electrons according to
the quantum mechanical model.
Section 5.2
Models of the Atom
 OBJECTIVES:
•
Describe how the shapes of
orbitals related to different
sublevels differ.
Ernest Rutherford’s Model

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Discovered dense positive piece
at the center of the atom“nucleus”
Electrons would surround and
move around it, like planets
around the sun
Atom is mostly empty space
It did not explain the chemical
properties of the elements – a
better description of the electron
behavior was needed
Niels Bohr’s Model
 Why
don’t the electrons fall into the
nucleus?
 Move like planets around the sun.
•
•
In specific circular paths, or orbits, at
different levels.
An amount of fixed energy separates
one level from another.
The Bohr Model of the Atom
I pictured the
electrons orbiting
the nucleus much
like planets
orbiting the sun.
Niels Bohr
However, electrons
are found in
specific circular
paths around the
nucleus, and can
jump from one
level to another.
Bohr’s model

Energy level of an electron
•
analogous to the rungs of a ladder
The electron cannot exist between
energy levels, just like you can’t stand
between rungs on a ladder
 A quantum of energy is the amount of
energy required to move an electron
from one energy level to another

The Quantum Mechanical
Model
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Energy is quantized - It comes in chunks.
A quantum 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.
In 1926, Erwin Schrodinger derived an
equation that described the energy and
position of the electrons in an atom
Schrodinger’s Wave Equation

d

V 
8  m dx
h
2
2
Erwin
Erwin Schrodinger
Schrodinger
2
2
 E
Equation for the
probability of a single
electron being found
along a single axis (x-axis)
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”
 An area where there
is a chance of
finding an electron.
 Think of fan blades

Atomic Orbitals
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Principal Quantum Number (n) = the
energy level of the electron: 1, 2, 3, etc.
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: letters s, p, d, and f
Principal Quantum Number
Generally symbolized by “n”, it denotes
the shell (energy level) in which the
electron is located.
Maximum number
of electrons that
can fit in an
energy level:
2n2
Summary
# of
Max
shapes electrons
Starts at
energy level
s
1
2
1
p
3
2
d
5
6
10
f
7
14
4
3
By Energy Level
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
First Energy Level
Has only s orbital
only 2 electrons
1s2
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Second Energy
Level
Has s and p orbitals
available
2 in s, 6 in p
2s22p6
8 total electrons
By Energy Level
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Third energy level
Has s, p, and d
orbitals
2 in s, 6 in p, and 10
in d
3s23p63d10
18 total electrons
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Fourth energy level
Has s, p, d, and f
orbitals
2 in s, 6 in p, 10 in
d, and 14 in f
4s24p64d104f14
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

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The orbitals do not
fill up in a neat
order.
The energy levels
overlap
Lowest energy fill
first.
Section 4.3
Electron Arrangement in Atoms

OBJECTIVES:
• Describe how to write the electron
configuration for an atom.
Section 4.3
Electron Arrangement in Atoms

OBJECTIVES:
• Explain why the actual electron
configurations for some elements
differ from those predicted by the
aufbau principle.
Using the Periodic Table to
Determine Electronic Arrangement
An electron configuration organizes an
atom’s electrons according to certain
rules. This configuration identifies the
energy in an atom.
 Electrons can only be in certain levels
according to the following pattern:
 1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,
5d,6p,7s,5f,6d,7p

BOHR MODELS OF
THE ATOM
Electrons by energy Level: (2n2)
 1st level
2 e max
 2nd level
8 e max
 3rd level
18 e max
 4th level
32 e max
 5th level
50 e max
 6th level
72 e max…
 7th level
98 e max

You will learn three ways to
write electron
configurations.
-Arrows
-Numbers
-Noble Gas Configuration
The Arrow Method
The Next couple of pages will teach you
how to complete the arrow method for
electron configuration.
 Remember, you must fill-up on letter
before moving on to a different type of
letter.
 Each space or dash can hold a maximum
of two arrows.
 Each space or dash (orbital) prefers to
be single unless it has to be paired up.

Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
Electron Configurations…
…are the way electrons are arranged in
various orbitals around the nuclei of
atoms. Three rules tell us how:
1) Aufbau principle - electrons enter the
lowest energy first.

•
This causes difficulties because of the
overlap of orbitals of different energies –
follow the diagram!
2) Pauli Exclusion Principle - at most 2
electrons per orbital - different spins
Pauli Exclusion Principle
No two electrons in an
atom can have the same
four quantum numbers.
Wolfgang Pauli
To show the different
direction of spin, a pair
in the same orbital is
written as:
Quantum Numbers
Each electron in an atom has a
unique set of 4 quantum numbers
which describe it.
1)
2)
3)
4)
Principal quantum number
Angular momentum quantum number
Magnetic quantum number
Spin quantum number
Electron Configurations
3) Hund’s Rule- When electrons
occupy orbitals of equal energy,
they don’t pair up until they have to.
Let’s write the electron configuration
for Phosphorus


We need to account for all 15
electrons in phosphorus
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3d
4s
3p
3s
2s
1s
The first two electrons
go into the 1s orbital
2p
Notice the opposite
direction of the spins
 only 13 more to go...

5f
4f
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
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p • The last three electrons
go into the 3p orbitals.
2p They each go into
separate shapes (Hund’s)
• 3 unpaired electrons
• = 1s22s22p63s23p3
Orbitals fill in an order
 Lowest
energy to higher energy.
 Adding electrons can change the
energy of the orbital. Full orbitals
are the absolute best situation.
 However, half filled orbitals have a
lower energy, and are next best
•
•
Makes them more stable.
Changes the filling order
Lets Try These:
Calcium
 Fluorine
 Titanium
 Copper
 Bromine

Numerical Electron
Configuration


Similar to the arrow method, but
instead of arrows, we use numerical
exponents.
Same rules apply.
Examples:


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
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

Hydrogen
Chlorine
Boron
Copper
Silicon
Calcium
Argon
Xenon
Silver
Zirconium


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





1s1
1s2 2s2 2p6 3s2 3p5
1s2 2s2 2p1
1s2 2s2 2p6 3s2 3p6 4s2 3d9
1s2 2s2 2p6 3s2 3p2
1s2 2s2 2p6 3s2 3p6 4s2
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
5p6
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d2
Noble Gas Electron
Configuration
(This is the shortest method)
**Color Periodic Table with s, p, d, f’s**
 You must indicate the last noble gas
element and then finish the
configuration.
 Calcium [Ar] 4s2

Examples:










Fluorine
Magnesium
Aluminum
Manganese
Tin
Iodine
Platinum
Cesium
Phosphorus
Nickel










[He] 2s2 2p5
[Ne] 3s2
[Ne] 3s2 3p1
[Ar] 4s2 3d5
[Kr] 5s2 4d10 5p2
[Kr] 5s2 4d10 5p5
[Xe] 6s2 4f14 5d8
[Xe] 6s1
[Ne] 3s2 3p3
[Ar] 4s2 3d8
Do These and Turn in:

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




Se
Y
Pd
Os
Au
Ra
Dy
Am
Ca
Mn
B









Ga
Mo
Te
Rb
Extra Credit (Long
Method)
V
In
Eu
Cr
Write the electron configurations for
these elements:



Titanium - 22
electrons
Vanadium - 23
electrons
Chromium - 24
electrons
1s22s22p63s23p64s23d2
2 2 6 2 6 2 3
 1s 2s 2p 3s 3p 4s 3d
2 2 6 2 6 2 4
 1s 2s 2p 3s 3p 4s 3d
(expected)
 But this is not what
happens!!

Chromium is actually:
 1s22s22p63s23p64s13d5
 Why?
 This
gives us two half filled orbitals
(the others are all still full)
 Half
full is 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 change gives one more filled
orbital and one that is half filled.


Remember these exceptions: d4,
d9
Irregular configurations of Cr and Cu
Chromium steals a 4s electron to
make its 3d sublevel HALF FULL
Copper steals a 4s electron
to FILL its 3d sublevel
Atomic Spectra
White light is
made up of all
the 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.
This is how we know
what stars are made
of.
• These are called
the atomic
emission
spectrum
• Unique to each
element, like
fingerprints!
• Very useful for
identifying
elements
Light is a Particle?
Energy is quantized.
 Light is a form of energy.
 Therefore, light must be quantized
 These smallest pieces of light are
called photons.
 Energy & frequency: directly related.

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, with only
one electron, and in the first energy
level.
Changing the energy

Heat, electricity, or light can move the
electron up to different energy levels. The
electron is now said to be “excited”
Changing the energy

As the electron falls back to the ground
state, it gives the energy back as light
Changing the energy


They may fall down in specific steps
Each step has a different energy
Ultraviolet
Visible
Infrared
The further they fall, more energy is
released and the higher the frequency.
 This is a simplified explanation!
 The orbitals also have different energies
inside energy levels
 All the electrons can move around.
