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Chapter 5
Electrons in Atoms
Greek Idea Democritus and
Leucippus Matter is made up
of indivisible particles
Dalton - one type of atom for each element
1 Summarize the development of the atomic theory in terms of the following:
Thompson - atom was a positive ball with electrons in it
Rutherford - atom had nucleus with electrons around it
Thomson’s Model
Discovered electrons Atoms were made of
positive stuff Negative electron
floating around “Plum-Pudding” model
Rutherford’s Model Discovered dense
positive piece at the center of the atom• Nucleus• Electrons
moved around• Mostly empty
space
Bohr’s Model 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.
Quantum mechanical model - atom has no definite shape
#1. Summarize the development of the atomic theory in terms of the following:
Bohr - atom had electrons in orbits around nucleus
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Bohr’s ModelIn
crea
sing
ene
rgy
Nucleus
First
Second
Third
Fourth
Fifth
} Further away
from the nucleus means more energy.
There is no “in between” energy
Energy Levels
#2. Define energy level and relate the energy level to the rungs of a ladder. Define quantum of energy.
There are energy levels around the nucleus. Think of a ladder
nucleus
1st
2nd
3rd
4th
These are the regions where an electron is likely to be found.
Electrons in levels farther from the nucleus contain more energy
Quantum of energy - the amount of energy that moves an electron up one level
Etc.
Waves vs. particles
Are electrons energy waves or particles? Both. They show properties of each and must be described as such: as particles - they have mass (1/1840 of a
proton) as energy waves - they can be emitted
from an atom as a measurable wave.
#3. Explain what is meant by the phrase “the energies of electrons are said to be quantized”.
Quantized energies means that each electron contains a specific amount of energy and that these amounts are not continuous, but units of energy. The electrons can’t “slide” from one energy level to another - they immediately appear in the next level..
#4. Explain the significance of quantized energies and the quantum mechanical model of the atom.
It is significant because it changes the way we look at energy. Instead of thinking that there are an infinite number of temperatures between 50ºC and 51ºC, we are forced to admit that there are a finite number of temperatures between these two degrees. And as something heats up, its temperature goes up by jumps rather than flowing smoothly.
#5. Explain why it is difficult to build a model of a quantum mechanical model of an atom
To build a model, you must know the where things are and their shape.
The electrons can’t be pinpointed and we only know where they are most likely to be found.
Sometimes electrons are even a long ways from the nucleus.
You can’t model anything this vague. Electrons also behave as waves and not
particles
The Bohr Model of the Atom
Neils Bohr
I pictured electrons orbiting the nucleus much like planets orbiting the sun.But I was wrong! They’re more like bees around a hive.
WRONG!!!
Quantum MechanicalModel of the Atom
Mathematical laws can identify the regions outside of the nucleus where electrons are most likely to be found.These laws are beyond the scope of this class. . . . .
But why? Hmmm lets see an example
Schrodinger Wave Equation
22
2 2
8dh EV
m dx
Equation for probabilityprobability of a single electron being found along a single axis (x-axis)
Erwin Schrodinger
Heisenberg Uncertainty Principle
You can find out where the electron is, but not where it is going.
OR…You can find out where the electron is going, but not where it is!
“One cannot simultaneously determine both the position and momentum of an electron.”
Werner Heisenberg
The Quantum Mechanical Model – Some Points
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.
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 – Some Points cont.
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 %
The Quantum Mechanical Model
#6. Distinguish among principal energy level, energy sublevel, and atomic orbital.
Principal energy level• Any of the major energy levels that contain
electrons - the rungs on the ladder• designated as 1, 2, 3, 4, etc.
Energy sublevel• Locations within the principal energy level
where specific electrons are found. Atomic orbital
• An orbital describes the shape of the energy sublevel and designated as s, p, d, & f orbitals
Old
Electron Energy Level (Shell)Generally symbolized by n, it denotes the probable distance of the electron from the nucleus. Number of electrons that can fit in a shell:
2n2
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.
An orbital is a region within an energy level where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…
1 energy sublevel: s orbital
2 energy sublevels:s & p orbitals
3 energy sublevels:s, p & d orbitals
4 energy sublevels:s, p, d & f orbitals
#7. State how the principal energy levels are designated and identify the number of sublevels within each principal energy level.
Old
nucleus
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.
1 s orbital for every energy level
Spherical shaped
Each s orbital can hold 2 electrons Called the 1s, 2s, 3s, etc.. orbitals.
S orbitals
Orbitals of the same shape (s, for instance) grow
larger as n increases…
Nodes are regions of low probability within an
orbital.
Sizes of s orbitals
P orbitals Start at the second energy level 3 different directions 3 different shapes Each can hold 2 electrons (opposite spins)
- three dumbbell-shaped p orbitals in
each energy level above n = 1,
- each assigned to its own axis (x, y and z) in space.
P orbital shape
P Orbitals - all together
D orbitals Start at the third energy level 5 different shapes Each can hold 2 electrons
Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of
“double dumbells”…and a “dumbell
with a donut”!
F orbitals Start at the fourth fourth energy level
Have seven different shapes
2 electrons per shape
Yes - There are no atoms big enough to require orbitals beyond F.
Summary
s
p
d
f
# of shapes
Max electrons
Starts at energy level
1 2 1
3 6 2
5 10 3
7 14 4
#7. State how the principal energy levels are designated and identify the number of sublevels within each principal energy level.
The principal energy levels are designated as n = 1, n = 2, n = 3, n = 4, n = 5, etc.
There may be as many as four sublevels (orbitals) in each principal energy level
1 has one sublevel
called “1s” with 1 orbital
2 has two sublevels
2s, 2p 2p has 3
orbitals
3 has three sublevels
3s, 3p, 3d 3d has 5
orbitals
4 has four sublevels
4s, 4p, 4d, 4f
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
The number of the principal energy level tells you how many sub energy levels (orbitals) are in that level.
The letters “ss”, “pp”, “dd” and “ff” tell you the shape of the orbital
By Energy Level First Energy Level only s orbital only 2 electrons 1s2
Second Energy Level s and p orbitals are
available 2 in s, 6 in p 2s22p6
8 total electrons
By Energy Level Third energy level s, p, and d orbitals 2 in s, 6 in p, and
10 in d 3s23p63d10
18 total electrons
Fourth energy level s,p,d, and f orbitals 2 in s, 6 in p, 10 in
d, ahd 14 in f 4s24p64d104f14
32 total electrons
Orbital filling table
5.2
Electron Configurations
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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
Writing electron configuration H = 1s1
He = 1s2
Li = 1s2 2s1
B = 1s2 2s2 2p1
Writing electron configuration C = 1s2 2s2 2p2
1s 2s 2p’s
N = 1s2 2s2 2p3
Write an electron configuration for fluorine
1s2 2s2 2p5
Write an electron configuration for boron
1s2 2s2 2p1
The first to electrons go into the 1s orbital
Notice the opposite spins
only 13 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
The next electrons go into the 2s orbital
only 11 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The next electrons go into the 2p orbital
• only 5 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The next electrons go into the 3s orbital
• only 3 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The last three electrons go into the 3p orbitals.
• They each go into seperate shapes
• 3 upaired electrons
• 1s22s22p63s23p3
The easy way to remember
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2
• 2 electrons
Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2
• 4 electrons
Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
• 12 electrons
Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2
• 20 electrons
Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6
5s2
• 38 electrons
Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6
5s2 4d10 5p6 6s2
• 56 electrons
Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6
5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
• 88 electrons
Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6
5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
5f14 6d10 7p6 • 108 electrons
The periodic table is arranged to show which orbital is being filled
s orbital p orbitald orbitalf orbital
State the maximum number of electrons that can go into each of the following sublevels:
2s
2 3p
6 4s
2 3d
10
4p
6 5s
2 4f
14 5p
6
State the number of electrons in the highest occupied energy level of these atoms.
Barium • 2p1 one
Sodium• 3s1 one
Aluminum• 3p1 one
Write electron configurations for atoms of these elements:
selenium• 1s2 2s2 2p6 3s2 3p6
4s2 3d10 4p4
• [Ar] 4s2 3d10 4p4
vanadium• 1s2 2s2 2p6 3s2 3p6
4s2 3d3
• [Ar] 4s2 3d3
nickel• 1s2 2s2 2p6 3s2 3p6
3s3 3p6 • [Ar] 4s2 3d8
calcium• 1s2 2s2 2p6 3s2 3p6
4s2
• [Ar] 4s2
oxygen• 1s2 2s2 2p4
• [He] 2s2 2p4
s orbital is being filleds orbital is
being filled
p orbital isbeing filled
d orbital is being filled
Exceptions to Electron Configuration
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 is 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
expect1s22s22p63s23p64s23d9
But the actual configuration is1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled orbital.
Remember these exceptions
Light/ WavesA Quick Review
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 kinds of waves
All move at 3.00 x 108 m/s ( c)
Parts of a wave
Wavelength
AmplitudeOrgin
Crest
Trough
Parts of Wave Origin - the base line of the energy. Crest - high point on a wave Trough - low point on a wave Amplitude - distance from origin to crest Wavelength - distance from crest to
crest• is abbreviated Greek letter lambda.
Frequency The number of waves that pass a
given point per second. Units are cycles/sec or hertz (hz) Abbreviated the Greek letter nu
c =
Frequency and wavelength Are inversely related As one goes up the other goes down. Different frequencies of light is
different colors of light. There is a wide variety of frequencies The whole range is called a spectrum
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low energy
High energy
Low Frequency
High Frequency
Long Wavelength
Short WavelengthVisible Light
Atomic SpectrumSec 5.3
How color tells us about atoms
Prism 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.
How we know what stars are made of.
Atoms can emit or absorb only specific frequencies of light
• These are called discontinuous spectra
• Or line spectra
• unique to each element.
• These are emission spectra
• The light is emitted given off.
Light is a Particle Energy is quantized. Light is energy Light must be quantized These smallest pieces of light are
called photons. Energy and frequency are directly
related.
Energy and frequency E = h x E is the energy of the photon is the frequency h is Planck’s constant
• h = 6.6262 x 10 -34 Joules sec. joule is the metric unit of Energy
The Math in Chapter 5 Only 2 equations
c = E = h
Plug and chug.
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?
An explanation of Atomic Spectra
So Where Does It Come From Anyway?
Where the electron starts When we write electron
configurations we are writing the lowest energy.
The energy level and electron starts from is called its ground state.
So……… Do the electrons ever change
orbitals?
YES!
For Electrons to Change Orbitals they
Must absorb energy to move up to a higher orbital (energy level)
Must release energy to move down to a lower orbital (energy level)
Energy in the Atom is Quantified
Max Plank theorized that electrons can absorb or emit only specific amounts of energy.
Quantum amounts – no more no less, must be exact.
Many Things are Quantified
Some are not
The Energy Involved is Light A photon
is a quantum of light energy.
Lets Watch What Happens http://spiff.rit.edu/classes/phys301/lec
tures/spec_lines/Atoms_Nav.swf
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
Changing the energy As the electron falls back to ground
state it gives the energy back as light
May fall down in steps Each with a different energy
Changing the 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.
Ultraviolet Visible Infrared
Emission Spectra of an Atom Acts like a fingerprint
Can Identify the type of atom from its spectral lines.
How helium was first discovered, by observing the Sun’s light spectra
Part of the Sun’s SpectrumDo you think the sun is made of more than one thing?
Why?
ALL the Light You SEE Was at one point emitted by an excited
atom.
So What is light? Light is a particle - it comes in chunks. Light is a wave - we can measure its
wave length and it behaves as a wave If we combine E=mc2 , c=, E = 1/2
mv2 and E = h We can get = h/mv The wavelength of a particle.
Matter is a Wave Does not apply to large objects Things bigger that an atom A baseball has a wavelength of about
10-32 m when moving 30 m/s An electron at the same speed has a
wavelength of 10-3 cm Big enough to measure.
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 because
Heisenberg Uncertainty Principle
It is impossible to know exactly the speed and velocity of a particle.
The better we know one, the less we know the other.
The act of measuring changes the properties.
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.
Moving Electron
Photon
Before
ElectronChanges velocity
Photon changes wavelength
After
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