Atomic Structure 1

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Atomic Structure
Bravo 15,000 kilotons
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Modern Atomic Theory

• As early as 400 B.C. Greek philosophers proposed
  that all matter is composed of atoms
• In 18th century Europe, the first chemists
  noticed characteristics shared by all compounds
• These observations of compounds and their
  reactions led to three important laws

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Law of Definite Proportions A given compound
contains the same elements in exactly the same
proportions by mass, regardless of the size or
source of the sample
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Law of Conservation of Mass In a reaction, matter
is neither created or destroyed Law of Multiple
Proportions When the same elements combine to
form different compounds, they do so in mass
ratios that can be expressed by small whole
numbers
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Daltons Atomic Theory (1803)
Many fine maidens are hot for me because of my
ruggedly handsome looks, boss sideburns, and my
atomic theory!
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Daltons Modern Atomic Theory

• All matter is made of indivisible and
  indestructible atoms.
• All atoms of a given element are identical in
  their chemical and physical properties
• Atoms of different elements differ in their
  chemical and physical properties.
• Atoms of different elements combine in simple
  whole number ratios to form compounds.
• Chemical reactions consist of the combination,
  separation, or rearrangement of atoms.

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Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to
deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
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Results of Cathode Ray Experiments

• Electrons travel in straight lines
• They have mass
• They are invisible
• They are independent of cathode composition
• They bend in a magnetic field like a
  negatively-charged particle would

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Thomsons Atomic Model
Got Plum Pudding? Yummy!
Thomson believed that the electrons were like
plums embedded in a positively charged pudding,
thus it was called the plum pudding model.
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Rutherfords Gold Foil Experiment

• Alpha particles are helium nuclei
• Particles were fired at a thin sheet of gold
  foil
• Particle hits on the detecting screen (film) are
  recorded

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Try it Yourself!
In the following pictures, there is a target
hidden by a cloud. To figure out the shape of the
target, we shot some beams into the cloud and
recorded where the beams came out. Can you figure
out the shape of the target?
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The Answers
Target 1
Target 2
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Rutherfords Findings

• Most of the particles passed right through
• A few particles were deflected
• VERY FEW were greatly deflected

Like howitzer shells bouncing off of tissue
paper! (I like Charmin the best!)
Rutherfords Conclusions

• The nucleus is small
• The nucleus is dense
• The nucleus is positively charged

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Atomic Particles
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The Atomic Scale

• Most of the mass of the atom is in the nucleus
  (protons and neutrons)
• Electrons are found outside of the nucleus (the
  electron cloud)
• Most of the volume of the atom is empty space

q is a particle called a quark
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About Quarks
Protons and neutrons are NOT fundamental
particles.
Protons are made of two up quarks and one
down quark.
Neutrons are made of one up quark and two
down quarks.
Quarks are held together by gluons
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Atomic Number
Atomic number (Z) of an element is the number of
protons in the nucleus of each atom of that
element.
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Mass Number
Mass number is the number of protons and neutrons
in the nucleus of an isotope.
Mass p n0
8
8
18
18
Arsenic
75
33
75
Phosphorus
15
31
16
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Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of
neutrons.
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Isotopes

• Atoms of an element with the same number of
  protons but different numbers of neutrons

• Most elements have more than one isotope

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Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Carbon 12.011
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Nuclear Symbols
Mass number (p no)
Element symbol
Atomic number (number of p)
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Cartoon courtesy of NearingZero.net
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In order to understand current atomic theory we
must first understand the properties of light
Light behaves both as a wave and as a particle
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Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the
electron as a particle.
His son, George Thomson won the Nobel prize for
describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!
Dont you get sassy with me, Boy!
Oh, go fly a kite, you old geezer!
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The Particle-like Electron
The photoelectric effect. Incoming EM radiation
on the left ejects electrons, depicted as flying
off to the right, from a substance. Only one
photon (a packet of light energy) can eject one
electron. Therefore, light acts like particles
(Einstein, 1905) because of the photons
quantized energy nature. This differed from
the description of EM radiation by Maxwell (1865)
which showed the infinite divisibility of EM
energy in physical systems.
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The Wave-like Electron
The electron propagates through space as an
energy wave. To understand the atom, one must
understand the behavior of electromagnetic waves.
Duuude! How you like my new doo? Aint it a trip?
Louis deBroglie
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Section 4.1
Light is a form of Electromagnetic Radiation It
exhibits wavelike behavior as it travel through
space Kinds of em radiation x-rays, ultraviolet
light, infared light, microwaves, radio
waves Together, all of the types of em radiation
form the electromagnetic spectrum
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Electromagnetic radiation propagates through
space as a wave moving at the speed of light.
c ??
C speed of light, a constant (3.00 x 108 m/s)
? frequency, in units of hertz (hz, sec-1)
? wavelength, in meters
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The energy (E ) of electromagnetic radiation is
directly proportional to the frequency (?) of the
radiation.
E h?
E Energy, in units of Joules (kgm2/s2)
h Plancks constant (6.626 x 10-34 Js)
? frequency, in units of hertz (hz, sec-1)
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Spectroscopic analysis of the visible spectrum
produces all of the colors in a continuous
spectrum
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Longer Wavelength, Lower Energy
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Types of electromagnetic radiation
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Wavelength Table
Long Wavelength Low Frequency Low ENERGY
Short Wavelength High Frequency High ENERGY
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Higher Frequency and Energy
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Electron transitionsinvolve jumps of definite
amounts ofenergy (quanta).
This produces bands of light with
definite wavelengths.
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Emission Spectrum
Continuous Emission Spectrum
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• Electron absorbs
• energy from the flame
• goes to a higher energy
• state.

2. Electron goes back down to lower energy state
and releases the energy it absorbed as light.
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lithium
sodium
potassium
copper
16.11
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Bohrs Model of the Atom (1913)

• e- can only have specific (quantized) energy
  values
• light is emitted as e- moves from one energy
  level to a lower energy level

n (principal quantum number) 1,2,3,
RH (Rydberg constant) 2.18 x 10-18J
7.3
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Schrodinger Wave Equation
Published in 1926, used the hypothesis that
electrons have a dual wave-particle nature
Together with the Heisenberg Uncertainty
Principle, it laid the foundation for modern
quantum theory Quantum Theory describes
mathematically the wave properties of electrons
and other very small particles
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The Bohr Model of the Atom
I pictured electrons orbiting the nucleus much
like planets orbiting the sun.
But I was wrong! Theyre more like bees around a
hive.
WRONG!!!
Neils Bohr
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Wave function equations only give the probability
of finding an electron at any given place around
the nucleus They do not travel around the nucleus
in neat orbits. Instead they exist in
orbitals Orbital A three-dimensional region
around the nucleus that indicates the probable
location of an electron.
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Quantum Numbers
Each electron in an atom has a unique set of 4
quantum numbers which describe it.

• Principal quantum number
• Angular momentum quantum number
• Magnetic quantum number
• Spin quantum number

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Think of the 4 quantum numbers as describing
where your seat is in the Superdome 1. Level 2.
Gate 3. Row 4. Seat
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Pauli Exclusion Principle
No two electrons in an atom can have the same
four quantum numbers.
And, no two fans in the Superdome should have the
same 4 seat numbers!
Wolfgang Pauli
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Principal Quantum Number
Generally symbolized by n, it denotes the shell
(energy level) in which the electron is located.
Number of electrons that can fit in a shell
2n2
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Angular Momentum Quantum Number
The angular momentum quantum number, generally
symbolized by l, denotes the orbital (subshell)
in which the electron is located.
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Magnetic Quantum Number
The magnetic quantum number, generally symbolized
by m, denotes the orientation of the electrons
orbital with respect to the three axes in space.
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Spin Quantum Number
Spin quantum number denotes the behavior
(direction of spin) of an electron within a
magnetic field.
Possibilities for electron spin
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Assigning the Numbers

• The three quantum numbers (n, l, and m) are
  integers.
• The principal quantum number (n) cannot be zero.
  
• n must be 1, 2, 3, etc.
• The angular momentum quantum number (l) can be
  any integer between 0 and n - 1.
• For n 3, l can be either 0, 1, or 2.
• The magnetic quantum number (m) can be any
  integer between -l and l.
• For l 2, m can be either -2, -1, 0, 1, or 2.

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Principle, angular momentum, and magnetic quantum
numbers n, l, and ml
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An orbital is a region within an atom where
thereis a probability of finding an electron.
This is a probability diagram for the s orbital
in the first energy level
Orbital shapes are defined as the surface that
contains 90 of the total electron probability.
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Sizes of s orbitals
Orbitals of the same shape (s, for instance) grow
larger as n increases
Nodes are regions of low probability within an
orbital.
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s orbital shape
The s orbital has a spherical shape centered
around the origin of the three axes in space.
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P orbital shape
There are three dumbbell-shaped p orbitals in
each energy level above n 1, each assigned to
its own axis (x, y and z) in space.
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d orbital shapes
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!
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Shape of the f orbitals
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• Max Planck proposed the idea of quanta, small
  specific amounts of energy.
• The electron cloud is the region outside of the
  nucleus where an electron can most probably be
  found.
• The Pauli exclusion principle states that no two
  electrons in the same atom can have the same four
  quantum numbers
• Hunds rule says that orbitals of equal energy are
  each occupied by one electron of the same spin
  before any is occupied by a second

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• Louis deBroglie believed that electrons could
  have a dual wave-particle nature
• The magnetic quantum number indicates the
  position of an orbital about the three axes in
  space.
• The photoelectric effect is the emission of
  electrons from metals that have absorbed photons.
• The wave model of light did not explain the
  photoelectric effect. Only the particle model
  could.
• The energy of a photon, or quantum, is related to
  its frequency.
• Both the Heisenberg uncertainty principle and the
  Schrodinger equation led to the concept of atomic
  orbitals

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Quantum Numbers (4) Specify the properties of
atomic orbitals and the properties of electrons
in orbitals Principal Quantum Number
(n) Indicates the main energy level occupied by
the electron Values are 1, 2, 3 etc
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Angular Momentum Quantum Number (l) Indicates the
shape of the orbital (values 0, 1, 2, or
3) (letter designation s, p, d, or f)
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Magnetic Quantum Number (m) Indicates the
orientation of an orbital around the nucleus
(values depend on orbital) for an s orbital, m
0 for a p orbital, m px, py, or pz (5 d
orientations, 7 f orientations)
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Spin Quantum Number Electrons can be thought of
as spinning on an internal axis. values are
either 1/2 or -1/2, which indicate the two
fundamental spin states of an electron
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Orbital filling table
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Principal Quantum Number
Generally symbolized by n, it denotes the shell
(energy level) in which the electron is located.
Number of electrons that can fit in a shell
2n2
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Angular Momentum Quantum Number
The angular momentum quantum number, generally
symbolized by l, denotes the orbital (subshell)
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Magnetic Quantum Number
The magnetic quantum number, generally symbolized
by m, denotes the orientation
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s orbital shape
The s orbital has a spherical shape centered
around the origin of the three axes in space.
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Sizes of s orbitals
Orbitals of the same shape (s, for instance) grow
larger as n increases
Nodes are regions of low probability within an
orbital.
74
P orbital shape
There are three dumbbell-shaped p orbitals in
each energy level above n 1, each assigned to
its own axis (x, y and z) in space.
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d orbital shapes
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!
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Shape of f orbitals
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Orbital Subshells
Subshell of orbitals e- in each
total e- s 1 2 2
p 3 2
6 d 5
2 10 f
7 2
14
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Electronic Structure of Atoms
Shells and Orbitals

• Shells of an atom contain a number of stacked
  orbitals

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Electronic Structure of Atoms
p
s
f
d
Relative Energies for Shells and Orbitals

• Some orbital subshells overlab others in
  different energy levels.

Relative Energies of the orbitals
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Orbital Filling Order
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Writing Electron Configurations

• Determine the number of electrons and atom has
• Fill orbitals in order of increasing energy

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Orbital Diagrams
Orbital Diagrams are models of electron
arrangements showing configuration, subshell,
aufbau, hunds, and pauli H ? 1S
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Orbital Diagrams
He ?? 1S Li ?? ?
1s 2s
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Hunds Rule
The most stable arrangement of electrons is that
with the most unpaired electrons all with the
same spin
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Practice Problems
1. Write the electron configuration and orbital
diagram for B
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Practice Problems

• Write the electron configuration and orbital
  diagram for B
• 1s2 2s2 2p1

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Practice Problems

• Write the electron configuration and orbital
  diagram for P

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Practice Problems

• Write the electron configuration and orbital
  diagram for P
• 1s2 2s2 2p6 3s2 3p3

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Practice Problems

• Write the electron configuration and orbital
  diagram for Sc

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Practice Problems

• Write the electron configuration and orbital
  diagram for Sc
• 1s2 2s2 2p6 3s2 3p6 4s2 3d1

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Practice Problems

• Write the electron configuration and orbital
  diagram for Pr

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Practice Problems

• Write the electron configuration and orbital
  diagram for Pr
• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
  4f3

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2.7 PT and Electron Configuration
The First 20 elements
1s

• H 1s1
• He 1s2
• Li 1s2 2s1
• Be 1s2 2s2
• B 1s2 2s2 2p1
• C 1s2 2s2 2p2
• N 1s2 2s2 2p3
• O 1s2 2s2 2p4
• F 1s2 2s2 2p5
• Ne 1s2 2s22p6
• Na 1s2 2s2 2p6 3s23s1
• Mg 1s2 2s2 2p6 3s2 3s2
• Al 1s2 2s2 2p6 3s2 3p1
• Si 1s2 2s2 2p6 3s2 3p2
• P 1s2 2s2 2p6 3 s 2 3p1
• S 1s2 2s2 2p6 3s2 3p4
• Cl 1s2 2s2 2p6 3s2 3p5
• Ar 1s2 2s2 2p6 3s2 3p6

2s
2p
3s
3p
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Orbital filling table
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Electronic Structure of Atoms
The First 10 elements
1s

• H 1s1
• He 1s2
• Li 1s2 2s1
• Be 1s2 2s2
• B 1s2 2s2 2p1
• C 1s2 2s2 2p2
• N 1s2 2s2 2p3
• O 1s2 2s2 2p4
• F 1s2 2s2 2p5
• Ne 1s2 2s2 2p6

2s
2p
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Pauli Exclusion Principle

• A maximum of two electrons can occupy each
  orbital. Each electron must have different spin
  quantum numbers

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Aufbau Principle
Electrons in an atom will occupy the
lowest-energy orbitals available (aufbau
building up)
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Hund's Rule
The most stable arrangement of electrons is that
with the maximum number of unpaired electrons,
all with the same spin quantum number