Modern Chemistry Chapter 4 Arrangement of Electrons in Atoms

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Modern ChemistryChapter 4Arrangement of
Electrons in Atoms
Sections 1-3 The Development of a New Atomic
Model The Quantum Model of the Atom Electron
Configurations
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2
Section 1
The Development of a New Atomic Model
Chapter 4 Section 1 New Atomic Model pages 97-103
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3
Electromagnetic radiation Electromagnetic
spectrum Wavelength Frequency Photoelectric
effect Quantum Photon Ground State Excited
state Line-emission spectrum Continuous
spectrum
Section 1 Vocabulary
Chapter 4 Section 1 New Atomic Model pages 97-103
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Properties of Light

• Electromagnetic Radiation a form of energy that
  exhibits wavelike behavior as it travels through
  space.
• Wavelength the distance between corresponding
  points on adjacent waves
• Frequency the number of waves that pass a given
  point in a specific time.

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p. 98
Electromagnetic Spectrum Image
Chapter 4 Section 1 New Atomic Model pages 97-103
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Properties of Light

• wavelength x frequency speed of light
• ? x ? c
• Visible Light R O Y G B I V

Long ? Low ? Low E
Short ? High ? High E
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The Photoelectric Effect

• The photoelectric effect refers to the emission
  of electrons from a metal when light shines on
  the metal.

p. 99
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The Photoelectric Effect

• Max Planck proposed that energy is proportional
  to the frequency of the electromagnetic wave.
• Electromagnetic energy is emitted from objects in
  small packages called quanta.
• E h ?
• Plancks constant
• h 6.626 x 10-34 Jsec

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The Photoelectric Effect

• Albert Einstein expanded on Plancks idea.
• Electromagnetic radiation has a dual
  wave-particle nature.
• A particle of light is a photon.
• Photon a particle of electromagnetic radiation
  having zero mass and carrying a quantum of
  energy.
• Ephoton h ?

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The Photoelectric Effect
To knock an electron loose, it must be hit with
a photon which possesses a minimum amount of
energy. This energy corresponds to its
frequency. Different metals hold electrons more
or less tightly So different metals
requiredifferent frequencies to show the
photoelectric effect.
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The H-Atoms Line Emission Spectrum

• Electric current is passed through a vacuum tube
  with hydrogen in it.
• A glow is produced
• When shined through a prism a line emission
  spectrum is produced

p. 101
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Emission Spectrum Animation
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Light and Energy

• The electromagnetic spectrum
• Particles (photons) and waves at the same time!
• c f?
• The greater f, the smaller ?
• The greater f, the higher the energy of the EM
  radiation.

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Which has higher energy red or blue light?
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• Emission spectra known for hundreds of years
• Niels Bohr unlocked their secret
• Electrons orbiting in shells around the nucleus
• Energy levels differ from one another

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Electrons and Light

• When an atom absorbs energy, electrons are
  promoted to higher energy levels.
• When the atom releases the absorbed energy, the
  electron falls back down to lower energy levels
  and EM radiation is given off (sometimes light)

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There are only certain posibilities

• Electrons can only exist at discrete energy
  levels
• Therefore, in changing levels, only certain
  amounts of energy can be released
• These amounts are unique to every element
• Page 343 in book illustrates

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How are the lines of the line spectrum formed?
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Spectrum
Excited State
n4
UV
Excited State
n3
Excited State unstable and drops back down
Vi s ible
Excited State
But only as far as n 2 this time
n2

• Energy released as a photon

• Frequency proportional to energy drop

IR
n1
Ground State
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Summary

• Electron normally in Ground State
• Energy supplied as heat or electricity
• Electron jumps to higher energy level
• Now in Excited State
• Unstable
• Drops back to a lower level

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Bohr Model of the Atom

• Niels Bohr proposed orbits for the electrons
• Each orbit has a fixed energy
• Lower energy orbits are closer to the nucleus
• Between orbits the electron cannot exist

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Explaining the Line Emission Spectrum




An electron absorbs a specific amount of energy
(absorption) and moves from its ground state to
an excited state
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Explaining the Line Emission Spectrum




The electron returns to its ground state and
emits a photon (emission).
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Explaining the Line Emission Spectrum




This photon has an energy corresponding to the
difference between the two states. This photon
has a specific E, ?, ? and color.
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Bohr Model of the Atom

• Bohrs calculated values for the lines agreed
  with the values observed for the lines in each
  series.
• However, it did not explain the spectraof atoms
  with more thanone electron.

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Bohr and Einstein

Chapter 4 Section 1 New Atomic Model pages 97-103
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Modern ChemistryChapter 4Arrangement of
Electrons in Atoms
Sections 1-3 The Development of a New Atomic
Model The Quantum Model of the Atom Electron
Configurations
28
29
Section 2
The Quantum Model of the Atom
Chapter 4 Section 2 The Quantum Model pages
104-110
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Heisenberg Uncertainty Principle Quantum theory
Orbital Quantum number Principal quantum
number Angular momentum quantum number Magnetic
quantum number Spin quantum number
Section 2 Vocabulary
Chapter 4 Section 2 The Quantum Model pages
104-110
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Electrons As Waves

• Electrons can have wave and particle
  characteristics (like light).
• Waves (electrons) confined to a space can have
  only certain frequencies
• The frequencies correspondto Bohrs orbits.
• Louis de Broglie

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Electrons As Waves

• This is confirmed by experiments
• Electrons, like waves, can be bent, diffracted
  and have interference
• Diffraction the bending of a wave as it passes
  through a smallopening
• Interference whenwaves overlap.

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The Heisenberg Uncertainty Principle

• Involves the detection of atoms
• To detect an electron a photon is used
• The photon interacts with the electron and
  changes its course
• There is uncertainty in trying to locate an
  electron.

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The Heisenberg Uncertainty Principle

• It is impossible to determine simultaneously both
  the position and velocity of an electron or any
  other particle.

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The Heisenberg Uncertainty Principle
Werner Heisenberg was stopped by the police.
When asked if he knew how fast he was going, he
replied
No officer, I dont know how fast I was going.
But I know exactly where I am.
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Heisenberg and Bohr

Chapter 4 Section 2 The Quantum Model pages
104-110
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The Schrödinger Wave Equation

• Developed an equation that treats electrons as
  waves
• Proved the quantization of electron energies
• Quantum theory describes mathematically the wave
  properties of electrons and other small
  particles

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The Schrödinger Wave Equation
Chapter 4 Section 2 The Quantum Model pages
104-110
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The Schrödinger Wave Equation

• Solutions to the equations are wave functions.
• Wave functions give the probability of finding
  electrons.
• Electrons do not travel in neat orbits.
• Electrons exist in regions called orbitals.
• Orbital a three dimensional region around the
  nucleus that indicates the probable location of
  electrons.

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Atomic Orbitals and Quantum Numbers

• Quantum numbers specify the properties of atomic
  orbitals and the properties of electrons in the
  orbital.
• The first three numbers result from the solution
  to Schrodingers wave equation.

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Principle Quantum Number
1

• Symbol n
• Energy Level
• n 1, 2, 3, 4, 5, 6, 7 (whole numbers)
• As n increases the electrons energy and average
  distance from the nucleus increases.

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Angular Momentum Quantum Number
2

• Symbol l
• Shape of orbital (sublevel)
• l s, p, d, f
• A sublevel is made up of a certain number of
  orbitals

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Atomic Orbitals and Quantum Numbers
One s orbital on the s sublevel
Chapter 4 Section 2 The Quantum Model pages
104-110
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Atomic Orbitals and Quantum Numbers
Three p orbitals on the p sublevel
Chapter 4 Section 2 The Quantum Model pages
104-110
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Atomic Orbitals and Quantum Numbers
Five d orbitals on the d sublevel
Chapter 4 Section 2 The Quantum Model pages
104-110
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Atomic Orbitals and Quantum Numbers
The 1st and 2nd energy levels orbitals
Chapter 4 Section 2 The Quantum Model pages
104-110
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7s 7p 7d 7f 7g 7h 7i 6s
6p 6d 6f 6g 6h 5s 5p 5d
5f 5g 4s 4p 4d 4f 3s 3p
3d 2s 2p 1s
of orbitals on a energy level
16 1s 3p 5d 7f
9 1 s 3p 5 d
4 1 s 3p
1 1 s only
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Atomic Orbitals and Quantum Numbers
Chapter 4 Section 2 The Quantum Model pages
104-110
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Magnetic Quantum Number
3

• Symbol m
• Orientation around the nucleus
• m x, y, z, xy, yz, xz

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Spin Quantum Number
4

• Symbol ms
• Spin State of the electron
• m 1/2, -1/2
• When electrons spin they produce a magnetic spin.
  
• Two electrons can exist in one orbital.
• Each electron must have an opposite spin state.

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7s2 7p6 7d10 7f14 7g 7h 7i 6s2
6p6 6d10 6f14 6g 6h 5s2 5p6 5d10
5f14 5g 4s2 4p6 4d10 4f14 3s2
3p6 3d10 2s2 2p6 1s2
of electronson a energy level
32e- 261014
18e- 2 6 10
8e- 2 6
2e- 2 in the s orbital
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7s2 7p6 7d10 7f14 7g 7h 7i 6s2
6p6 6d10 6f14 6g 6h 5s2 5p6 5d10
5f14 5g 4s2 4p6 4d10 4f14 3s2
3p6 3d10 2s2 2p6 1s2
of electronson a energy level
32e- 261014
18e- 2 6 10
8e- 2 6
2e- 2 in the s orbital
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Atomic Orbitals and Quantum Numbers
Chapter 4 Section 2 The Quantum Model pages
104-110
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Modern ChemistryChapter 4Arrangement of
Electrons in Atoms

• Sections 1-3
• The Development of a New Atomic Model
• The Quantum Model of the Atom
• Electron Configurations

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Section 3
Electron Configurations
Chapter 4 Section 3 Electron Configurations pages
110-122
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Electron configuration Aufbau principle Pauli
exclusion principle Hunds Rule Noble Gas
Noble Gas configurations
Section 3 Vocabulary
Chapter 4 Section 3 Electron Configurations pages
110-122
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Rules Governing Electron Configurations

• Aufbau Principle
• An electron occupies the lowest energy orbital
  that can receive it
• Less energy is required for electrons to pair up
  in the 4s than the 3d.
• Aufbau German for construction

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7s 7p 7d 7f 7g 7h 7i 6s
6p 6d 6f 6g 6h 5s 5p 5d
5f 5g 4s 4p 4d 4f 3s 3p
3d 2s 2p 1s
Sublevel Filling Order
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Rules Governing Electron Configurations

• Pauli Exclusion Principle
• No two electrons in the same atom can have the
  same set of four quantum numbers.
• Electrons must have opposite spin states
• Two electrons can exist in an orbital.

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Rules Governing Electron Configurations

• Hunds Rule
• Orbitals of equal energy are each occupied by one
  electron before any orbital is occupied by a
  second electron, andall electrons in singly
  occupied orbitals must have the same spin state.
• Repulsion between electrons is minimized.

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Representing Electron Configurations

• Orbital Notation
• ?? ?? ? ? ? . 1s 2s 2p
• Which element is this?
• How does this show Hunds rule?

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Orbital Diagram for Si
orbital
14 electrons
electron (-1/2 spin)
1s
2s
2p
3s
3p
electron (1/2 spin)
sublevel
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Shapes of s, p and d orbitalsImage
p.108
Chapter 4 Section 3 Electron Configurations pages
110-122
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Representing Electron Configurations




Orbital Notation ?? ?? ? ? ? . 1s
2s 2p

• Electron Configurations
• 1s2 2s2 2p3
• Use superscripts instead of lines and arrows.

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Electron Configurations
For Manganese 25 electrons 1s 2s 2p 3s
3p 4s 3d
6
2
2
6
2
2
5
number of electrons in the sublevel
sublevel
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PRACTICE

• Page 113 1 2
• Write the orbital diagram for Carbon
• Write the orbital diagram for Beryllium
• Write the electron configuration for P.
• Write the electron configuration for Cl.

p. 113
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Elements of the Second Period




3Li 1s2 2s2 6C 1s2 2s2 2p2
Highest occupied energy level C 3 valence
electronsLi 1 valence electron
Inner shell electrons
Valence electrons electrons occupying the
highest energy level in an atom
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Elements of the Third Period
10Ne 11Na 12Mg 18Ar 19K
1s2 2s2 2p6 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2
1s2 2s2 2p6 3s2 3p6 1s2 2s2 2p6 3s2 3p6 4s1
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Elements of the Third Period
Look! Its neon!
10Ne 11Na 12Mg 18Ar 19K
1s2 2s2 2p6 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2
1s2 2s2 2p6 3s2 3p6 1s2 2s2 2p6 3s2 3p6 4s1
Ar?!
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Elements of the Third Period
10Ne 11Na 12Mg 18Ar 19K
1s2 2s2 2p6 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2
1s2 2s2 2p6 3s2 3p6 1s2 2s2 2p6 3s2 3p6 4s1
Ne
Ne
Ar
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Elements of the Third Period
10Ne 11Na 12Mg 18Ar 19K
1s2 2s2 2p6 Ne 3s1 Ne 3s2 1s2 2s2 2p6 3s2
3p6 Ar 4s1
Chapter 4 Section 3 Electron Configurations pages
110-122
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Noble Gas Configuration

• Find the noble gas with an atomic number closest
  to but less than the elements atomic number.
• Find the next sublevel after that noble gas
• Fill in sublevels with the leftover electrons.
  (atomic of element atomic of noble gas)

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Noble Gas Configurations
1s2 2p6 3p6 4p6 5p6 6p6
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Chapter 4 Section 3 Electron Configurations pages
110-122
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1s
2s
2p
3s
3p
4s
3d
4p
4d
5p
5s
6s
5d
6p
7s
6d
4f
5f
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Elements of the Fourth Period

• Deviations
• 24Cr
• 24Cr Ar 3d5 4s1
• unpaired electrons givea more stable arrangement
  witha lower energy
• 29Cu
• 29Cu Ar 3d10 4s1
• No explanation for either

WRONG!!!
Ar 3d6
Ar 3d9 4s2
WRONG!!!
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Elements of the Fourth Period

• Even though the d sublevel fills before the p
  sublevel, the s sublevel is moved to be with the
  p sublevel.
• 53I Kr 5s2 4d10 5p5 is written as
• 53I Kr 4d10 5s2 5p5
• The sublevels on the same energy level are
  together.
• It also shows the valence electrons.

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Elements of the Fifth Period




Y Kr 4d1 5s2Zr Kr 4d2 5s2 Nb Kr
4d3 5s2 Mo Kr 4d4 5s2 Tc Kr 4d5 5s2
Ru Kr 4d6 5s2Rh Kr 4d7 5s2 Pd
Kr 4d8 5s2 Ag Kr 4d9 5s2 Cd Kr
4d10 5s2
Nb Kr 4d4 5s1 Mo Kr 4d5 5s1 Tc
Kr 4d6 5s1 Ru Kr 4d7 5s1Rh Kr
4d8 5s1 Pd Kr 4d10 Ag Kr 4d10 5s1
p.120
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PRACTICE
Page 121 1-4
p. 113
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Elements of the Sixth Period

• 4f and 5d are very close in energycausing many
  deviations
• Look at the configurations on the periodic table
  on the back cover of the book.

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PRACTICE
Page 122 1 2
p. 113
1. answer 2. answer 3. answer
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Section 3 Homework
Chapter 4 Section 3 Electron Configurations pages
110-122
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