Download Unit 4 - School District of Durand

10 Modern Atomic Theory and the
Periodic Table
The amazing colors of fireworks result from electron transfer
between energy levels of atoms.
Foundations of College Chemistry, 14th Ed.
Morris Hein and Susan Arena
Copyright © 2014 John Wiley & Sons, Inc. All rights reserved.
Learning Objectives
Electromagnetic Radiation
• List the 3 basic characteristics of electromagnetic radiationwavelength, frequency and speed
• Calculate wavelength, frequency using the speed of light.
The Bohr Atom
• Explain the relationship between the line spectrum and the
quantized energy levels of the electron cloud.
• Describe the Lyman, Paschen and Balmer series exhibited in the
hydrogen model as discovered by Bohr .
• Differentiate between continuous, emission and absorption
spectrums.
• Describe the history of quantum theory.
Modern Atomic Model-Quantum Atom Model
• Describe the principal energy levels, sublevels and orbitals of
an atom through quantum numbers n, l.
• Explain the relationship between energy and electron position.
Atomic Structure
• Use quantum numbers to find elements on the periodic table.
• Use the guidelines to write full, short-cut and box diagram
electron configurations.
• Describe orbitals of an atom with quantum numbers
ml , ms .
• Describe and electron using 4 quantum number.
Electron Structures and the Periodic Table
• Describe how the electron configurations of the atoms relate to
their position and chemical reactivity.
• Predict Oxidation states using electron configuration.
Electromagnetic Radiation
Electromagnetic Radiation:
A form of energy that runs a continuum from radio to
X-rays, visible light to microwaves.
Each form of radiation shares common characteristics:
they display wavelike properties and travel at the same speed.
Basic Properties of Waves
Wavelength (λ): The distance between two similar points in
consecutive waves; such as from crest to crest or trough to
trough. SI Units = meters, m
Frequency (n or f ): The number of waves that pass a point per
unit of time. SI Units = Hertz, Hz or 1/sec=(s-1)
Wavelength and frequency are inversely related.
Speed (v or s): how fast a wave moves through space.
Parts and Properties of Waves
Amplitude (Ψ ):The distance between the highest or lowest
point on a wave to the resting line. SI Units = meters, m
Crest:: The highest point in a wave.
Trough: The lowest point in a wave.
Resting line: The imaginary line that travels through the
middle of the wave. Ambassador
Waves transmit Energy!!! The amount of energy present is
seen in the amplitude and frequency of the wave.
Electromagnetic Radiation
Electromagnetic Radiation
Electromagnetic Spectrum:
The full range of electromagnetic radiation,
arranged based on wavelength.
Electromagnetic radiation has both wave-like
and particle properties.
Radiation can behave like tiny packets (“bundles”)
of energy called photons.
© 2014 John Wiley & Sons, Inc. All rights reserved.
Speed of Light “C” = frequency x wavelength
c = lf
c = 3.00 x 108 m/s
3.00 x 108 m/s =
lf
1. a) Calculate the wavelength of a radio wave with a frequency of
1.1x106 Hz.
f =1.1x106 Hz
3.00 x 108 m/s = l (1.1x106 Hz)
3.00 x 108 m/s =
270 m
1.1x106 Hz
b)
What part of the EMF does this fall into?
Radio snd TV waves.
How many waves will pass in a 5.0 seconds?
Recall, frequency describes how many waves pass per second, thus
1.1x106Hz = 1.1x106waves per second
1.1x106 waves x 5.0 seconds = 5.5x106 waves.
second
2. What is the frequency of an electromagnetic wave that has a
wavelength of 2.56x10-12m?
3.00 x 108 m/s = l f
3.00 x 108 m/s = 2.56x10-12m∙f
3.00 x 108 m/s =f = 1.17x1020 Hz
2.56x10-12m
What EMF wave has a
wavelength of 2.56x10-12m?
A gamma ray or very
energetic x-ray.
The Bohr Atom
At high temperatures or when high voltages are applied,
elements radiate (emit) colored light.
When this light is passed through a prism,
a set of brightly colored lines result.
Line spectrum of hydrogen
These line spectra indicate the light emitted has only
specific wavelengths/frequencies.
Each element possesses a characteristic
and unique line spectrum.
© 2014 John Wiley & Sons, Inc. All rights reserved.
The Bohr Atom
How can these experimental results be explained?
From his study of the line spectrum of hydrogen,
Bohr proposed a revised theory of the atom.
Bohr suggested electrons exist in
specific regions at defined
distances from the nucleus.
The electrons then move about
the nucleus in circular orbits at a
fixed distance from the nucleus.
© 2014 John Wiley & Sons, Inc. All rights reserved.
Bohr Atom Continued
Bohr’s Model of the Hydrogen Atom
FYI: 95% of all atoms in the universe are hydrogen.
Hydrogen atoms exist in only specified energy states.
Hydrogen atoms can absorb only certain amount of energy, and no
others.
3. When excited hydrogen atoms lose energy, they lose only certain
amount of energy, emitted as photons.
4. The different photons given off by hydrogen atoms produce the color
line seen in the bright-light spectrum of hydrogen.
i)The greater the energy lost by the atom, the greater the
energy of the photon.
1.
2.
The Bohr Atom Continued…
Bohr also suggested energy absorbed or emitted
by an atom is quantized (has discrete fixed units).
Bohr proposed that electrons can orbit
the nucleus at different distances.
Each orbit is a distinct, discrete (quantized) energy level.
When an atom absorbs energy,
the electrons can be promoted to
higher energy levels.
When an atom emits energy,
the electrons can decays to
a lower energy level.
© 2014 John Wiley & Sons, Inc. All rights reserved.
The Bohr Atom
Ground state: lowest energy level for an atom.
Each line in the spectrum corresponds to emission
of energy as an electron relaxes from a
higher to lower energy level.
Color of light emitted depends on the gap
between the energy levels.
Bohr’s theory worked very well to explain and predict
the line spectrum of hydrogen…. BUT….
Bohr’s theory broke down in multi-electron systems.
© 2014 John Wiley & Sons, Inc. All rights reserved.
n = 4 ionization
n = 3 second excited state
Paschen (IR)
n = 2 first excited state
Balmer (Visible)
n = 1 Ground state
Lyman (UV)
Lyman Series: Produces Ultraviolet light waves
Excited electrons drop from a high energy level back to
energy level 1 to produce UV light.
Balmer Series: Produces visible light in the form of color
Excited electrons drop from a high energy level back to
energy level 2 to produce visible light.
Paschen Series: Produces infrared light.
Excited electrons drop from a high energy level back to
energy level 3 to produce infrared light.
Spectrum Facts
• Electrons will remain at the lowest possible energy level that produces
the most atomic stability.
• More energy is required to exist farther away from the nucleus. This
equates to higher potential energy.
• Electrons that absorb energy can move to an energy level further away
from the nucleus.
• When the electrons return to the lower energy level, they give back the
energy, in the form of light (color), that was absorbed.
Go Packers
Atomic Spectra
•
•
•
Discrete energy levels in any atom or molecule
(quantum mechanics)
Transitions between levels produces (or absorbs) quanta
of light, also called photons.
Bigger energy steps produce more powerful photons
(ie..shorter wavelength, blue, UV)
Emission Spectra’s show discrete colored lines. The lines correspond to
photons of discrete energies that are emitted when excited atomic states in
the gas make transitions back to lower levels. Background in black.
A continuum spectrum is an emission spectrum where the lines over-lap
with each other and we an no longer distinguish the individual emission
lines.
An absorption spectrum occurs when light passes through a cold, dilute gas
and atoms in the gas absorb energy at specific frequencies. The absorption
spectrum is the opposite of an emission spectrum.
Continuous Spectrum
Emission Spectrum
Absorption Spectrum
I believe
electrons
move like
waves.
Albert Einstein
Get a grip
Plank,
electrons move
like particles
Max Plank
•1901Plank found that atoms can
only adsorb and emit energy in
distinct quantities; (Quanta) this
showed that energy displayed
particle-like properties.
Planck’s equation for electron
energy is E = hf
h = 6.6264x10-31 J-s
f = frequency
•1905: Einstein
suggested that energy itself is
quantized and can be viewed as a
string of particles called photons.
He established that energy has
mass. Einstein described electron
motion through the famous
equation E = mc2
m = mass
c= speed of light
A photon checks into a hotel and is asked if he needs any
help with his luggage.
He says, "No, I'm traveling light."
Ha Ha Ha!!!!!!!!!!!!!!!!!!!!!!!!
Photoelectric effect
•Einstein used Planck’s equation to explain the photoelectric effect
•Electrons are ejected from the surface of a metal when light shines on the
metal.
•For each metal, a minimum frequency of light is needed to release electrons.
• Red light has too low a frequency to cause ejection of electrons from
sodium while faint violet light releases electrons easily.
Einstein proposed that light consists of quanta of
Energy that behave like tiny particles of light.
He called these energy quanta photons.
Einstein suggests every photon carries an amount
of energy described by Planck’s equation: E = hn
deBroglie (Mid-1920’s) found a way to please
them both: the “Wave-particle duality of Nature”
Einstein
E=
mc2
Planck
and hf = E
mc2 = hf
mc2 = h c
l
Since both equations equal
energy, deBroglie set them
equal to each other.
Recall, c = lf
Thus, f = c/ l
Substitute…
The final equation
l = hc = h
Solve equation for l
simplify
predicts the wavelength
mc2
mc
of a particle at a given
l = h
Let c = velocity
mass and velocity. Thus
mv
combining both men’s ideas.
Ultimately, this suggests electrons act both like waves and particles.
Both matter and radiation possess a
remarkable duality of character, as they
sometimes exhibit the properties of waves, at
other times those of particles. Now it is
obvious that a thing cannot be a form of wave
motion and composed of particles at the same
time - the two concepts are too different.
(Werner Heisenberg, on Quantum Theory,
1930)
The famous Heisenberg uncertainty principle was first proposed by
Werner Heisenberg in 1927. According to this principle, it is
impossible to simultaneously measure the position and momentum of
a particle (exactly). Indeed, a good knowledge of the particle's
position implies a poor knowledge of its momentum, and vice versa.
Note that the uncertainty principle is a direct
consequence of representing particles as waves.
In the mid-1920s, Erwin Schrodinger, building on
the dual nature of matter, began focusing on the
wave-like properties of the electron by visualizing
them like standing waves.
Schrodinger treatment of the electron as a wave
enabled him to develop a mathematical and
graphical model that described electron behavior.
His mathematical model gave us the
position and shape of electron orbitals.
He suggests that 90% of the time,
electrons with a designated amount
of energy existed in a specific orbital.
Basically, he figured out the
shapes for s, p, d and f orbitals.
Max Born verified
Schrodinger’s probability.
Arthur Compton (1892-1962) demonstrated that a photon could collide with an electron.
Incident
Photon
Motion of
Electron after
collision
(Random
energy)
Motion of photon
After collision
The Darkroom Mystery
You have a friend who is taking a photography class at school. She is learning to be a good photographer as
well as how to process her film. So she regularly uses a darkroom for developing her picture. The darkroom is
equipped with a dim red light so that your friend can see enough to handle the film and the developing solution.
One day when the red bulb burned out, your friend replaced it with a dim yellow bulb that she found. She and
several other students were later dismayed when they ruined their film while using the darkroom. When your
friend told you of this disaster, you immediately understood what had happened. Explain what happened to
your friend’s film.
Arthur Compton verified energy has
mass, as Einstein suggested. He
demonstrated that a photon could collide
with an electron.
In 1925 Wolfgang Pauli enunciated his
exclusion principle - "under no
circumstances whatsoever may two
electrons have precisely the same set of
quantum numbers"
The Dual Particle/Wave Nature of Matter
Quantum mechanics deal in electron probabilities;
orbits from Bohr theory are replaced by orbitals.
Orbitals: regions of space with a high probability of
finding an electron.
An orbital for a hydrogen atom.
© 2014 John Wiley & Sons, Inc. All rights reserved.
Modern Atomic Model  Quantum Theory
Electron Cloud ( Home of the electrons.)
· Region, surrounding the nucleus, made up of
mostly empty space.
· Can be described as the atoms volume.
· A region where 90% of the time you will find
electrons is called an orbital.
Electrons
· Only moveable part of the atom.
· Housed in specific regions corresponding to
energy requirements; orbitals.
· Electrons are described using four-quantum
numbers.
n, l, ml , ms
n
l
n = Principal quantum number
ml
ms
 Describes the energy level of the electron which
often corresponds to the row number
 n = 1, 2, 3, 4, etc..
 Maximum number of electrons in a principal quantum
number is 2n2
o when n = 2; 2 (2)2 = 8, when n = 4; 2(4)2= 32
 Each energy level is divided into sublevels.
o The sublevels have different energy
requirements.
 The number of sublevels in each principal energy
level corresponds to the principal quantum number.
o Example. When n = 2, that means there are
2 sublevels in this principal energy level.
Energy Levels of Electrons
Bohr’s idea of quantized energy levels does have
parallels in quantum mechanics.
For example, quantum mechanics predicts discrete,
quantized principal energy levels for electrons in an atom.
Principal energy level (n):
provides a general idea of the distance
of an electron from the nucleus
(as n increases, so does the distance
from the nucleus)
n can be a positive integer (1,2,3, etc.)
© 2014 John Wiley & Sons, Inc. All rights reserved.
l=
Sublevel Shape
 There are four sublevel shapes and they are
described by the letters s, p, d, and f.
 Each of the sublevels are composed of orbitals.
 Orbitals are areas that can hold a maximum of 2
electrons.
o s = sharp (composed of 1 orbital)
Quantum number = 0
o p = principal (composed of 3 orbitals)
Quantum number = 1
o d = diffuse (composed of 5 orbitals)
Quantum number = 2
o f = fundamental (composed of 7 orbitals)
Quantum number = 3
1s < 2s < 3s < 4s < 5s < 6s < 7s
P-sublevel looks like dumbells
The p-sublevel continues to get further away from
The nucleus as “n” increases.
Principal quantum level n =3
Contains three sublevels/orbitals (3s, 3p and 3d orbitals)
The 5 3d orbitals have unique shapes relative
to s and p orbitals.
10 total electrons can occupy the
5 d orbitals of a subshell.
© 2014 John Wiley & Sons, Inc. All rights reserved.
Principal quantum level n =4
Contains four sublevels/orbitals -(4s, 4p, 4d and 4f )
The 7 - 4f orbitals have unique shapes relative
to s and p orbitals.
14 total electrons can occupy the 7 f orbitals.
As the electron cloud build, the region fills with the uniquely shaped orbitals.
Electron Configurations are AWESOME!!!!
Electron configurations:
•Show the number of electrons in the atom
•Show which principal quantum levels are involved
•Show which sublevels are involved
•More information can be gleaned from the electron
configuration and thus…there is more information to come.
1s22s22p63s23p63d104s24p64d104f145s25p65d16s2
Luteium
The Aufbau principle is really a thought process in which we think
about building up an atom from the one that precedes it in atomic
number, by adding a proton and neutrons to the nucleus and one
electron to the appropriate atomic orbital.
Recall…
The number of sublevels ,
s,p,d,f, in each principal
energy level, 1, 2, 3…7,
corresponds to the
principal quantum number.
This holds true through
energy level 4.
When
n =1 
•Energy level 5 only uses
4 sublevels,
•Energy level 6 only uses
3 sublevels
•Energy level 7 only uses
2 sublevels.
6
2
3
4
5
7
1s
2s 2p
3s 3p
4s 4p
5s 5p
6s 6p
7s 7p
3d
4d 4f
5d 5f
6d
We simply do not have enough elements to require more than 4
sublevels at this time.
Electrons enter the lowest
energy sublevels first. The
energy requirements of some
sublevels may surprise you.
Let’s take a look.
Notice that sublevel 4s is
filled before sublevel 3d.
That is due to the lower
energy requirement for 4s
when compared to 3d.
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
4f
5f
As you followed the arrows you
saw the pathway for entering the lowest energy levels first.
The s and p sublevels have row numbers and principal
quantum numbers that correspond. The row numbers and
principal quantum numbers of the d and f sublevels do not
correspond.
Electron Filling
Sublevel Filling Diagram
© 2014 John Wiley & Sons, Inc. All rights reserved.
Three types of electron configurations: Lutetium
Full Configuration:
1s22s22p63s23p63d104s24p64d104f145s25p65d16s2
Short-cut configuration:
[Xe] 4f145d16s2
Box Diagram Configuration Yikes… you need to show every
electron arrow! So let’s NOT do lutetium!!
Sodium :
Full:
Box :
1s22s22p63s1
The box diagram uses the full configuration and it represents each
electron with an arrow.
Short-cut:
1s2
[Ne] 3s1
2s2
2p6
3s1
ml = Magnetic orientation
 This letter describes the orbital’s magnetic orientation
about the x, y, and z axes.
 ml tells us how the electron cloud surrounding the
nucleus is directed in space.
 ml is described numerically using 0, 1, 2, 3, 4
and so on.
o The + sign indicates a positive direction, the – sign
indicates a negative direction.
o ml for the s orbital= 0
o ml for the p orbitals= 0, 1 (-1, 0, 1)
o ml for the d orbitals= 0, 1, 2 (-2, -1, 0, 1, 2)
o ml for the f orbitals= 0, 1,2,3, (-3, -2,-1, 0, 1, 2, 3)
-3
-2
-1
0
1
2
3
How many electrons can fit in an orbital?
Pauli Exclusion Principle: an atomic orbital can hold two
electrons, which must have opposite spins.
Electron spin is represented by arrows (  or  )
ms
= Electron Spin (±½)
 This quantum number describes the direction of the
electron’s spin.
 The 2 electrons in each orbital will spin in opposite
directions to reduce repulsion.
Each electron can be described by assigning the 4 quantum numbers
3d7
n
3
l
2
s=0, p=1
d=2, f=3
Recall there are 5d orbitals
-2
ml
-1
-1
ms
-½
0
1
Electrons enter
1 at a time.
±2, ±1, 0Numbers > 5 are -½
Numbers < 5 are +½
2
Level
n
1
2
3
4
5
6
7
Total number of
electrons in level, 2n2
2(1)2 = 2
2(2)2 = 8
2(3)2 = 18
2(4)2 = 32
2(5)2 = 50
2(6)2 = 72
2(7)2 =98
Number of
sublevels, n
1: s
2: s, p
3: s, p, d
4: s, p, d, f
5: s, p, d, f, ?
6: s, p, d, f, ?, ?
7: s, p, d, f, ?, ?, ?
Sublevel
Orbitals, n2
(1)2= 1
(2)2= 4
(3)2= 9
(4)2= 16
(5)2= 25
(6)2= 36
(7)2= 49
Energy Levels of Electrons Practice
What is the maximum number of electrons that can
occupy the n = 3 sublevel?
a. 8
b. 2
c. 18
d. 10
The n = 3 sublevel has 3 types of orbitals:
s (1), p (3), and d (5).
Two electrons can occupy each orbital.
(9 x 2 = 18)
© 2014 John Wiley & Sons, Inc. All rights reserved.
Atomic Structure of the First 18 Elements
Valence Electrons:
electrons located in the highest energy
(outermost) orbitals of an atom.
Oxygen
Example
Electron configuration
1s22s22p4
Outermost valence electrons are in the n =2 subshell;
oxygen has 6 total valence electrons.
The column number (1A-7A) in the periodic table
gives the valence electrons of an element.
Valence electrons participate in bonding to form
molecular compounds (see Chapter 11).
© 2014 John Wiley & Sons, Inc. All rights reserved.
Predicting Oxidation State
• We can predict the oxidation state of elements by looking at all the valence
electrons relative to stability.
• Transition and inner transition metals can have more than one oxidation state
because now that some electrons are housed in the “d or f” orbitals.
• There are different ways to enhance stability.
Rules for Stability:
 Most stable: all sublevels are full
 2nd most stable: all sublevels are full or half-full
 3rd most stable: d or f sublevels enlist promotion of 1 or 2 electrons from the

s-orbital to create a half-full d or f sublevel.
 Least stable-desirable: one sublevel is neither full or half-full
o d and f sublevels can have oxidation states result using promotion (columns 6
and 11) or one or more oxidation states resulting from giving away different
numbers of electrons. Learn by doing!
Predicting Oxidation States of Copper
1 electron from 4s2 can be promoted to 3d10
1s22s22p63s23p64s23d9
1s22s22p63s23p64s13d10
Written numerically, it is easy to see what electron
will be given away during chemical reactions.
1s22s22p63s23p63d10 4s1
= Cu+1
If no promotion occurs, what electrons will be given away?
Resulting in
1s22s22p63s23p63d9
= Cu+2
Using the short-Cut form is even easier
[Ar] 4s23d9
[Ar] 4s23d9
[Ar] 4s13d10 = Cu+1
= Cu+2
[Ar] 3d9
4s2
Electron Filling Practice
The electron configuration [Ar] 4s1 is the
ground state electron configuration of:
a. K
b. P
c. Fluorine
The element contains 1 valence electron
in the fourth period (due to n = 4).
d. Na
© 2014 John Wiley & Sons, Inc. All rights reserved.
Electron Filling Practice
The electron configuration [Ne] 3s23p1, is the ground
state electron configuration of:
a. Na
b. Al
c. Ar
The element contains 3 valence electrons
in the third period (due to n = 3).
d. S
© 2014 John Wiley & Sons, Inc. All rights reserved.
Chemical Periodicity
Groups of elements show similar chemical properties
because of similarities in their valence electrons.
Group number equals the total number of
outermost valence electrons.
Example Group 7 – contains the
ns2np5 valence configuration
© 2014 John Wiley & Sons, Inc. All rights reserved.