Download Atomic Structure and Periodicity Part 1

ATOMIC STRUCTURE
AND PERIODICITY
PART 1
CHAPTER 7
WARM-UP QUESTION
Be prepared to share out your response to the
following questions.
What is a photon?
What is the source of electromagnetic waves?
Is the color spectrum simply a small segment of the
electromagnetic spectrum? Defend your answer.
ELECTROMAGNETIC
RADIATION
WARM-UP CONTINUED
What is a photon?
• A particle of light.
• Particle vs Wave Theory
• Video 1
• Video 2
What is the source of electromagnetic waves?
• Accelerating electric charges
Is the color spectrum simply a small segment of the
electromagnetic spectrum? Defend your answer.
• Yes; the spectrum is also made up of radio waves, IR, UV, Xrays, and gamma rays.
CHARACTERISTICS OF WAVES
Waves are described according
to their
Amplitude
measures DISPLACEMENT
size of the disturbance (from
rest to crest)
Wavelength 
distance of a “repeating unit”
Also called a cycle
Velocity v
speed = how fast wave travels
FREQUENCY V

How often
number of wavelengths that pass any point per second
measured in
wavelengths/second or
cycles/second
Hertz (Hz) = number of
wavelengths in 1 second
Frequency is related to
velocity: c = v 
ELECTROMAGNETIC WAVE
a transverse wave with an
electric component and a
magnetic component at right
angles to each other
How are electromagnetic
waves (ex: light) different
from mechanical waves
(ex: sound and slinky)?
micro.magnet.fsu.edu
ELECTROMAGNETIC
WAVES
Electromagnetic waves are special
in the fact that they do not need a
medium to propagate through.
But what is creating the
disturbance? What is emitting this
energy?
© 2003 Mike Maloney
8
ELECTRONS
ELECTROMAGNETIC
WAVES
Electrons in materials are vibrated and emit energy in
the form of photons, which propagate across the
universe.
Photons have no mass, but are pure energy.
Electromagnetic Waves are waves that are made up of
these “photons”.
© 2003 Mike Maloney
9
When these photons come in contact with boundaries,
E-M waves interact like other waves would.
© 2003 Mike Maloney
10
ELECTROMAGNETIC SPECTRUM
SPEED OF E/M WAVES
It has been found that the speed of E-M waves and light
is ---
© 2003 Mike Maloney
11
• 3 x 108 or 300,000,000 m/s
• 671,000,000 mph
• 186,000 miles per second
• We call this value “c”
C is constant throughout the universe,
as long as light is in a vacuum.
When it is in other materials, c can
change, but can never be larger than its
value in a vacuum.
© 2003 Mike Maloney
12
Since “c” is constant, all of E-M waves
will have a corresponding frequency to
go along with their wavelength.
ENERGY IN E-M
WAVES
Which waves have more energy, Radio
waves or gamma waves?
The greater the frequency of an E-M
wave, the more crests pass a point in a
certain amount of time, therefore the
more photons pass that point.
© 2003 Mike Maloney
13
This means that more energy moves
past that point in a certain amount of
time or that the wave contains more
energy.
ELECTROMAGNETIC SPECTRUM
“CHECK-UP”
True or False…
1. Blue light has a shorter wavelength than red
light.
2. X-rays have lower frequencies than radio
waves.
3. Microwaves have higher frequencies than
gamma rays.
4. Visible radiation composes the major portion of
the electromagnetic spectrum.
True; False; False; False
WAVELENGTH-FREQUENCY
RELATIONSHIP EXAMPLE
Photosynthesis uses light with a frequency of
4.54x1014s-1. What wavelength does this
correspond to?
A: 660nm
WAVELENGTH-FREQUENCY
RELATIONSHIP PRACTICE
Calculate the frequency of blue light of wavelength
4.5 x 102nm.
Calculate the wavelength of green light of frequency
5.7 x 1014Hz.
A:6.7x1014Hz ; 5.3 x 10-7m or 530nm
THE NATURE OF MATTER
ΔE = hv = hc/λ
• ΔE is the change in energy for a system (in Joules per
photon)
• h is Planck’s constant (6.626 x 10-34J s)
• experimentally determined
• v is the frequency of the wave (s-1 or Hz)
**Energy can be gained or lost only in integer multiples
of hv. (quanta)
ENERGY, FREQUENCY,
WAVELENGTH EXAMPLE
Sodium atoms have a characteristic yellow color
when excited in a flame. The color comes from the
emission of 589.0nm.
• What is the frequency of this radiation?
• What is the change in energy associated with
this photon? Per mole of photons?
ENERGY, FREQUENCY,
WAVELENGTH PRACTICE
It takes 382 kJ of energy to remove one mole of
electrons from gaseous cesium. What is the
wavelength associated with this energy?
Would we be able to “directly” observe this energy
change? Why or why not.
THE PHOTOELECTRIC EFFECT
Emission of electrons from a metal when light shines on the
metal
Electromagnetic radiation (light) strikes the surface of the metal
ejecting electrons from the metal and causing an electric
current, if the frequency was below a certain minimum.
Analysis of the kinetic energy and numbers of the emitted
electrons led Einstein to suggest that electromagnetic radiation
can be viewed as a stream of photons.
*Note that the apparent mass of a photon depends on its
wavelength. The mass of a photon at rest is thought to be zero,
although we never observe it at rest.*
BIG IDEAS FROM EINSTEIN
AND PLANCK
• Energy is quantized. It can occur only in discrete
units called quanta.
• Electromagnetic radiation, which was previously
thought to exhibit only wave properties, seems
to show certain characteristics of particulate
matter as well. (dual nature of light)
WAVE-LIKE BEHAVIOR
Diffraction
• Light is scattered from a regular array of points
or lines.
• Constructive interference
• In-phase (bright)
• Destructive interference
• Out-of phase (dim/dark)
ATOMIC SPECTRUM OF HYDROGEN
Continuous Spectrum
• Contains all the wavelengths over which the
spectrum is continuous
Line Spectrum
• Contains certain specific wavelengths that are
characteristic of the substance emitting those
wavelengths
*Hydrogen’s line spectrum shoes that only certain
energy transfers are allowed in hydrogen.
*Specific energy levels among which the hydrogen
electron can shift, thus energy levels are quantized.
INDIVIDUAL PRACTICE
15, 31, 33, 35, 39
THE BOHR MODEL
1913 Niels Bohr developed the Quantum Model for the hydrogen
atom.
• The electron in the hydrogen atom moves around the nucleus
only in certain allowed circular orbits.
• Hydrogen atom energy levels consistent with the hydrogen
emission spectrum. (different wavelength/color associated with
the different levels of emission)
• Ground state
• The lowest possible energy state of an atom or molecule
• Excited state
• Higher potential energy state than ground state of an atom or
molecule
Although Bohr’s model fits the energy levels for
hydrogen, it is a fundamentally incorrect model for
the hydrogen atom.
Bohr’s model paved the way for later theories on
the quantization of energy in atoms.
Electrons do NOT move around the nucleus in
circular orbits (planetary model).
QUANTUM MECHANICS
de Broglie and Schrodinger – wavelike properties of
electrons
A specific wave function (function of the coordinates x, y,
and z of the electron’s position in 3-D space) is often called
an orbital.
• The wave function corresponding to the lowest energy for
the hydrogen atom is called the 1s orbital (no association
to the Bohr “orbit”).
Nature of an orbital takes into consideration the work of
Heisenberg.
• Heisenberg uncertainty principle: There is a fundamental
limitation to just how precisely we can know both the
position and momentum of a particle at a given time.
1S ORBITAL
The definition most often used by chemists to describe the
size of the hydrogen 1s orbital is the radius of the sphere that
encloses 90% of the total electron probability.
(90% of the time the electron I in this sphere)
SUMMARY
In the quantum (wave) mechanical model, the electron is viewed as a
standing wave. This representation leads to a series of wave
functions (orbitals) that describe the possible energies and spatial
distributions available to the electron.
In agreement with the Heisenberg uncertainty principle, the model
cannot specify the detailed electron motions. Instead, the square of
the wave function represents the probability distribution of the
electron in that orbital. This allows us to picture orbitals in terms of
probability distributions, or electron density maps.
The size of an orbital is arbitrarily defined as the surface that
contains 90% of the total electron probability.
The hydrogen atom has many types of orbitals. In the ground state,
the single electrons resides in the 1s orbital. The electron can be
excited to higher-energy orbitals if energy is put into the atom.
QUANTUM MECHANICS
HTTP://WWW.META-SYNTHESIS.COM/WEBBOOK/30_TIMELINE/310PX-BOHRATOM-PAR.SVG.PNG
Better than any previous model,
quantum mechanics does explain how
the atom behaves.
Quantum mechanics treats electrons
not as particles, but more as waves
(like light waves) which can gain or
lose energy.
But they can’t gain or lose just any
amount of energy. They gain or lose a
“quantum” of energy.
A quantum is just an amount of energy that the electron
needs to gain (or lose) to move to the next energy level.
In this case it is losing the energy and dropping a level.
ATOMIC ORBITALS
HTTP://MILESMATHIS.COM/BOHR2.JPG
Much like the Bohr model, the energy
levels in quantum mechanics describe
locations where you are likely to find
an electron.
Remember that orbitals are “geometric
shapes” around the nucleus where
electrons are found.
Quantum mechanics calculates the
probabilities where you are “likely” to
find electrons.
ATOMIC ORBITALS
HTTP://COURSES.CHEM.PSU.EDU/CHEM210/QUANTUM/QUANTUM.HTML
Of course, you could find an electron anywhere if
you looked hard enough.
So scientists agreed to limit these calculations to
locations where there was at least a 90% chance of
finding an electron.
Think of orbitals as sort of a "border” for spaces
around the nucleus inside which electrons are
allowed. No more than 2 electrons can ever be in 1
orbital. The orbital just defines an “area” where you
can find an electron.
What is the chance of finding an electron in the
nucleus? Yes, of course, it’s zero. There are not any
electrons in the nucleus.
ENERGY LEVELS
HTTP://WWW.CHEM4KIDS.COM/FILES/ART/ELEM_PERTABLE2.GIF
Quantum mechanics has a
principal quantum number. It is
represented by a little n. It
represents the “energy level”
similar to Bohr’s model.
Red
Orange
Yellow
Green
Blue
Indigo
Violet
n=1
n=2
n=3
n=4
n=5
n=6
n=7
• n=1 describes the first energy level
• n=2 describes the second energy
level
• Etc.
Each energy level represents a
period or row on the periodic table.
It’s amazing how all this stuff just
“fits” together.
SUB-LEVELS = SPECIFIC
ATOMIC ORBITALS
Each energy level has 1 or more
“sub-levels” which describe the
specific “atomic orbitals” for that
level.
Blue = s block
• n = 1 has 1 sub-level (the “s” orbital)
• n = 2 has 2 sub-levels (“s” and “p”)
• n = 3 has 3 sub-levels (“s”, “p” and
“d”)
• n = 4 has 4 sub-levels (“s”, “p”, “d”
and “f”)
There are 4 types of atomic orbitals:
• s, p, d and f
• Each of these sub-levels represent the
blocks on the periodic table.
ORBITALS
HTTP://MEDIA-2.WEB.BRITANNICA.COM/EB-MEDIA/54/3254-004-AEC1FB42.GIF
HTTP://UPLOAD.WIKIMEDIA.ORG/WIKIPEDIA/COMMONS/THUMB/E/E1/D_O
RBITALS.SVG/744PX-D_ORBITALS.SVG.PNG
s
p
d
In the s block, electrons are going into s orbitals.
In the p block, the s orbitals are full. New electrons are going into the p orbitals.
In the d block, the s and p orbitals are full. New electrons are going into the d orbitals.
What about the f block?
QUANTUM NUMBERS
Describe the properties of the orbital.
Name
Symbol
Property of the Orbital
Related to size and energy of the
orbital
Range of
Values
Principal
Quantum
Number
n
Integers
Angular
Momentum
Quantum
Number
l
Related to the shape of the orbital
Integers from
“subshell”
n-1 to 0
0 is s; 1 is p; 2 is d; 3 is f; 4 is g; 5 is h
Magnetic
Quantum
Number
ml
Related to the position of the orbital
in space relative to other orbitals
1 to ∞
Integers from
-l to 0 to +l
DEGENRATE
All orbitals having the same value of “n” have the
same energy.
3s; 3p; 3d
Energy is required to transfer an electron to a
higher-energy orbital (excited state).
**In polyelectronic atoms we find that the s, p, and
d have different levels of potential energy.
THE 4TH QUANTUM NUMBER
The electron spin quantum number.
Electron spin
• Two spin states + ½ and – ½
• Produce two oppositely directed magnetic moments
Pauli Exclusion Principle
• In a given atom no two electrons can have the same set of
four quantum numbers (n, l, ml, ms)
• Thus, an orbital can hold only TWO electrons, and they must
have opposite spins.
PRACTICE WITH QUANTUM
NUMBERS
Which of the following quantum numbers are
allowed? For each that is incorrect state why.
Principal, Angular Momentum, Magnetic
Quantum Numbers (n, l, ml)
a. 1, 0, 1
b. 2, 2, 1
c. 5, 3, 2
d. 6, -2, 2
e. 6, 2, -2
QUANTUM NUMBERS
AND LEVELS OF
ORBITALS
Table 7.2 on page 294 in text
Energy
Level
Sublevels
Total Orbitals
Total
Electrons
Total Electrons
per Level
n=1
s
1 (1s orbital)
2
2
n=2
s
p
1 (2s orbital)
3 (2p orbitals)
2
6
8
n=3
s
p
d
1 (3s orbital)
3 (3p orbitals)
5 (3d orbitals)
2
6
10
18
Complete the chart in your notes as we discuss this.
first
has an s orbital.
n = The
4
s level (n=1)
1 (4s orbital)
2 It has only
32 1.
There pare no other
energy level.
3 (4p orbitals
orbitals) in the first
6
d
5 (4d orbitals)
10
We call
this
orbital
the
1s
orbital.
f
7 (4f orbitals)
14
WHERE ARE THESE
ORBITALS?
HTTP://WWW.BIOSULF.ORG/1/IMAGES/PERIODICTABLE.PNG
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
4f
5f
ELECTRON
CONFIGURATIONS
What do I mean by “electron
configuration?”
The electron configuration is the
specific way in which the atomic
orbitals are filled.
Think of it as being similar to your
address. The electron configuration
tells me where all the electrons “live.”
RULES FOR ELECTON CONFIGURATIONS
HTTPS://TEACH.LANECC.EDU/GAUDIAS/SCHEME.GIF
In order to write an electron
configuration, we need to know the
RULES.
3 rules govern electron
configurations.
• Aufbau Principle
• Pauli Exclusion Principle
• Hund’s Rule
Using the orbital filling diagram at
the right will help you figure out
HOW to write them
• Start with the 1s orbital. Fill each
orbital completely and then go to the
next one, until all of the elements have
been acounted for.
Each line represents
an orbital.
1 (s), 3 (p), 5 (d), 7 (f)
High Energy
FILL LOWER ENERGY
ORBITALS FIRST
HTTP://WWW.META-SYNTHESIS.COM/WEBBOOK/34_QN/QN3.JPG
The Aufbau Principle states
that electrons enter the
lowest energy orbitals first.
The lower the principal
quantum number (n) the
lower the energy.
Within an energy level, s
orbitals are the lowest
energy, followed by p, d and
then f. F orbitals are the
highest energy for that level.
Low Energy
NO MORE THAN 2 ELECTRONS
IN ANY ORBITAL…EVER.
HTTP://WWW.FNAL.GOV/PUB/INQUIRING/TIMELINE/IMAGES/PAULI.JPG
Wolfgang Pauli, yet
another German
Nobel Prize winner
The next rule is the Pauli Exclusion Principal.
The Pauli Exclusion Principle states that an
atomic orbital may have up to 2 electrons and
then it is full.
The spins have to be paired.
We usually represent this with an up arrow and
a down arrow.
Since there is only 1 s orbital per energy level,
only 2 electrons fill that orbital.
Quantum numbers describe an electrons position, and no 2
electrons can have the exact same quantum numbers. Because of
that, electrons must have opposite spins from each other in order
to “share” the same orbital.
HUND’S RULE
HTTP://INTRO.CHEM.OKSTATE.EDU/AP/2004NORMAN/CHAPTER7/LEC111000.HTML
Don’t pair up the 2p electrons
until all 3 orbitals are half full.
Hunds Rule states that when you
get to degenerate orbitals, you fill
them all half way first, and then
you start pairing up the electrons.
What are degenerate orbitals?
Degenerate means they have the
same energy.
So, the 3 p orbitals on each level
are degenerate, because they all
have the same energy.
Similarly, the d and f orbitals are
degenerate too.
Paramagnetic
unpaired electrons
2p
Diamagnetic
all electrons paired
2p
APPLICATION
NOW that we know the rules, we can try to write some electron
configurations.
Remember to use your orbital filling guide/PERIODIC TABLE
to determine WHICH orbital comes next.
Lets write some electron configurations for the first few
elements, and let’s start with hydrogen.
H; Li; B; N; F; Na; K; Fe
ELECTRON
CONFIGURATIONS
Element
Configuration
Element
Configuration
H Z=1
1s1
He Z=2
1s2
Li Z=3
1s22s1
Be Z=4
1s22s2
B
Z=5
1s22s22p1
C
Z=6
1s22s22p2
N Z=7
1s22s22p3
O
Z=8
1s22s22p4
F
1s22s22p5
Ne Z=10
1s22s22p6
(2p is now full)
Na Z=11
1s22s22p63s1
Cl Z=17
1s22s22p63s23p5
K Z=19
1s22s22p63s23p64s1
Sc Z=21
1s22s22p63s23p64s23d1
Fe Z=26
1s22s22p63s23p64s23d6
Br Z=35
1s22s22p63s23p64s23d104p5
Z=9
Note that all the numbers in the electron configuration add up to the atomic
number for that element. Ex: for Ne (Z=10), 2+2+6 = 10
CONCEPTUAL CHECK
One last thing. Look at the previous slide and look
at just hydrogen, lithium, sodium and potassium.
Notice their electron configurations. Do you see any
similarities?
Since H and Li and Na and K are all in Group 1A,
they all have a similar ending. (s1)
ELECTRON
CONFIGURATIONS
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K Z=19
1s22s22p63s23p64s1
This similar configuration causes them to behave the
same chemically.
It’s for that reason they are in the same family or group
on the periodic table.
Each group will have the same ending configuration, in
this case something that ends in s1.
NOBLE GAS NOTATION…
“SHORT CUT”
Be
Al
Br
Mo
Ag
ORBITAL NOTATION
Be
Al
N
Br
Mo
Ag
ION ELECTRON
CONFIGURATION
Be2+
Al3+
Br-
Ag1+
INDIVIDUAL PRACTICE
Bond Energies- page 384 (on Unit 3 Test!)
• # 53 and 54
Periodicity and Atomic Structure-starting on page
322
• # 57, 58, 59, 60, 62, 67, 69, 70, 71, 72, 73, 74, 85,
87, 89, 95, 97, and 123
(Please note that there are multiple questions over the same
concept(s). You do not need to do them all but need to KNOW
how to do them.)