Electronic Structure of Atoms

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Electronic Structure of Atoms

• Resources
• Our TB Ch. 6 of Chemistry The Central Science
  AP version (10th edition)
• Powerpoint (from pearson) and in-class work
• POGIL activities (1) Analysis of Spectral Lines
  and (2) Interaction of Radiation and Matter
• Online resources for our TB (in particular online
  quiz)
• Chem tours from ch. 7 of the W.W. Norton online
  book by Gilbert
• http//www.wwnorton.com/college/chemistry/chemistr
  y3/ch/07/studyplan.aspx

• Animations from Glencoe site
• http//glencoe.mcgraw-hill.com/sites/0023654666/s
  tudent_view0/chapter7/
• Extra quizzes from Glencoe
• http//glencoe.mcgraw-hill.com/sites/0023654666/s
  tudent_view0/chapter7/
• Video lectures from chem guy
• http//www.kentchemistry.com/moviesfiles/chemguy/
  AP/ChemguyAtomicTheory.htm
• Handouts and practice problems from M. Brophys
  web site

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Chapter 6Electronic Structureof Atoms
Chemistry, The Central Science, 10th
edition Theodore L. Brown H. Eugene LeMay, Jr.
and Bruce E. Bursten
John D. Bookstaver St. Charles Community
College St. Peters, MO ? 2006, Prentice Hall, Inc.
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Waves

• To understand the electronic structure of atoms,
  one must understand the nature of electromagnetic
  radiation.
• The distance between corresponding points on
  adjacent waves is the wavelength (?).

4
Waves

• The number of waves passing a given point per
  unit of time is the frequency (?).
• For waves traveling at the same velocity, the
  longer the wavelength, the smaller the frequency.

5
Electromagnetic Radiation

• All electromagnetic radiation travels at the same
  velocity the speed of light (c), 3.00 ? 108
  m/s.
• Therefore,
• c ??

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The Nature of Energy

• The wave nature of light does not explain how an
  object can glow when its temperature increases.
• Max Planck explained it by assuming that energy
  comes in packets called quanta.

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The Nature of Energy

• Einstein used this assumption to explain the
  photoelectric effect.
• He concluded that energy is proportional to
  frequency
• E h?
• where h is Plancks constant, 6.63 ? 10-34 J-s
  (i.e. units for h are Js)

8
The Nature of Energy

• Therefore, if one knows the wavelength of light,
  one can calculate the energy in one photon, or
  packet, of that light
• c ??
• E h?

9

• For electromagnetic radiation animation and
  problems see
• http//www.wwnorton.com/college/chemistry/gilbert
  2/tutorials/interface.asp?chapterchapter_07folde
  rfrequency_wavelength

• For All Chem tours for the electrons in atoms and
  periodic properties topic see
• http//www.wwnorton.com/college/chemistry/chemistr
  y3/ch/07/studyplan.aspx
• Recommeded chem tours animations
• Electromagnetic radiation
• Light Emission and Absorbtion
• Bohr Model of the Atom
• De Broglie Wavelngth
• Quantum numbers
• Electron configuration

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The Nature of Energy

• Another mystery involved the emission spectra
  observed from energy emitted by atoms and
  molecules.

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The Nature of Energy

• One does not observe a continuous spectrum, as
  one gets from a white light source.
• Only a line spectrum of discrete wavelengths is
  observed.

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Go To Glencoe Animation

• http//glencoe.com/sites/common_assets/advanced_pl
  acement/chemistry_chang9e/animations/chang_7e_esp/
  pem1s3_1.swf
• POGIL activity on Spectral Lines
• (To Complete)

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The Nature of Energy

• Niels Bohr adopted Plancks assumption and
  explained these phenomena in this way
• Electrons in an atom can only occupy certain
  orbits (corresponding to certain energies).

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The Nature of Energy

• Niels Bohr adopted Plancks assumption and
  explained these phenomena in this way
• Electrons in permitted orbits have specific,
  allowed energies these energies will not be
  radiated from the atom.

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The Nature of Energy

• Niels Bohr adopted Plancks assumption and
  explained these phenomena in this way
• Energy is only absorbed or emitted in such a way
  as to move an electron from one allowed energy
  state to another the energy is defined by
• E h?

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The Nature of Energy

• The energy absorbed or emitted from the process
  of electron promotion or demotion can be
  calculated by the equation

where RH is the Rydberg constant, 2.18 ? 10-18 J,
and ni and nf are the initial and final energy
levels of the electron.
17
Go To Glencoe and Norton Animations

• http//glencoe.com/sites/common_assets/advanced_pl
  acement/chemistry_chang9e/animations/chang_7e_esp/
  pem1s3_1.swf
• POGIL activity on Interaction of Radiation and
  Matter
• (To Complete)
• Go to Chem tour for Bohr Model of atom (and
  Rydberg equation)
• http//www.wwnorton.com/college/chemistry/chemistr
  y3/ch/07/chemtours.aspx

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The Wave Nature of Matter

• Louis de Broglie posited that if light can have
  material properties, matter should exhibit wave
  properties.
• He demonstrated that the relationship between
  mass and wavelength was

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The Uncertainty Principle

• Heisenberg showed that the more precisely the
  momentum of a particle is known, the less
  precisely is its position known
• In many cases, our uncertainty of the whereabouts
  of an electron is greater than the size of the
  atom itself!

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Quantum Mechanics

• Erwin Schrödinger developed a mathematical
  treatment into which both the wave and particle
  nature of matter could be incorporated.
• It is known as quantum mechanics.

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The Quantum Mechanical Model

• Energy is quantized - It comes in chunks.
• A quantum 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.
• In 1926, Erwin Schrodinger derived an equation
  that described the energy and position of the
  electrons in an atom
• (this slide from J. Hushens presentation on
  Atomic Structure at http//teachers.greenville.k12
  .sc.us/sites/jhushen/Pages/AP20Chemistry.aspx)

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Schrodingers Wave Equation
Equation for the probability of a single
electron being found along a single axis (x-axis)
Erwin Schrodinger
(this slide from J. Hushens presentation on
Atomic Structure at http//teachers.greenville.k12
.sc.us/sites/jhushen/Pages/AP20Chemistry.aspx)
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Quantum Mechanics

• The wave equation is designated with a lower case
  Greek psi (?).
• The square of the wave equation, ?2, gives a
  probability density map of where an electron has
  a certain statistical likelihood of being at any
  given instant in time.

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Quantum Numbers

• Solving the wave equation gives a set of wave
  functions, or orbitals, and their corresponding
  energies.
• Each orbital describes a spatial distribution of
  electron density.
• An orbital is described by a set of three quantum
  numbers.

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Principal Quantum Number, n

• The principal quantum number, n, describes the
  energy level on which the orbital resides.
• The values of n are integers 0.

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Azimuthal Quantum Number, l

• This quantum number defines the shape of the
  orbital.
• Allowed values of l are integers ranging from 0
  to n - 1.
• We use letter designations to communicate the
  different values of l and, therefore, the shapes
  and types of orbitals.

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Azimuthal Quantum Number, l
Value of l 0 1 2 3
Type of orbital s p d f
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Magnetic Quantum Number, ml

• Describes the three-dimensional orientation of
  the orbital.
• Values are integers ranging from -l to l
• -l ml l.
• Therefore, on any given energy level, there can
  be up to 1 s orbital, 3 p orbitals, 5 d orbitals,
  7 f orbitals, etc.

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Magnetic Quantum Number, ml

• Orbitals with the same value of n form a shell.
• Different orbital types within a shell are
  subshells.

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s Orbitals

• Value of l 0.
• Spherical in shape.
• Radius of sphere increases with increasing value
  of n.

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s Orbitals

• Observing a graph of probabilities of finding an
  electron versus distance from the nucleus, we see
  that s orbitals possess n-1 nodes, or regions
  where there is 0 probability of finding an
  electron.

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p Orbitals

• Value of l 1.
• Have two lobes with a node between them.

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d Orbitals

• Value of l is 2.
• Four of the five orbitals have 4 lobes the other
  resembles a p orbital with a doughnut around the
  center.

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Energies of Orbitals

• For a one-electron hydrogen atom, orbitals on the
  same energy level have the same energy.
• That is, they are degenerate.

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Energies of Orbitals

• As the number of electrons increases, though, so
  does the repulsion between them.
• Therefore, in many-electron atoms, orbitals on
  the same energy level are no longer degenerate.

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Spin Quantum Number, ms

• In the 1920s, it was discovered that two
  electrons in the same orbital do not have exactly
  the same energy.
• The spin of an electron describes its magnetic
  field, which affects its energy.

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Spin Quantum Number, ms

• This led to a fourth quantum number, the spin
  quantum number, ms.
• The spin quantum number has only 2 allowed
  values 1/2 and -1/2.

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Pauli Exclusion Principle

• No two electrons in the same atom can have
  exactly the same energy.
• For example, no two electrons in the same atom
  can have identical sets of quantum numbers.

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Go To www.ptable.com

• IMPORTANT
• Use periodic Table to help you write electron
  configurations of atoms (and ions)

• Dynamic Periodic Table and Investigate (play
  with) the Orbitals option (on Top Tabs) for
  quantum numbers, orbitals and electron
  configurations of various elements
• Go To
• Glencoe site for animations on electron
  configuration
• http//glencoe.mcgraw-hill.com/sites/0023654666/st
  udent_view0/chapter7/animations_center.html

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Electron Configurations

• Distribution of all electrons in an atom
• Consist of
• Number denoting the energy level

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Electron Configurations

• Distribution of all electrons in an atom
• Consist of
• Number denoting the energy level
• Letter denoting the type of orbital

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Electron Configurations

• Distribution of all electrons in an atom.
• Consist of
• Number denoting the energy level.
• Letter denoting the type of orbital.
• Superscript denoting the number of electrons in
  those orbitals.

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Orbital Diagrams

• Each box represents one orbital.
• Half-arrows represent the electrons.
• The direction of the arrow represents the spin of
  the electron.

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Hunds Rule

• For degenerate orbitals, the lowest energy is
  attained when the number of electrons with the
  same spin is maximized.

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Periodic Table

• We fill orbitals in increasing order of energy.
• Different blocks on the periodic table, then
  correspond to different types of orbitals.

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Some Anomalies

• Some irregularities occur when there are enough
  electrons to half-fill s and d orbitals on a
  given row.

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Some Anomalies

• For instance, the electron configuration for
  chromium is
• Ar 4s1 3d5
• rather than the expected
• Ar 4s2 3d4.
• N.B. Copper is another
• anomaly.

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Some Anomalies

• This occurs because the 4s and 3d orbitals are
  very close in energy.
• These anomalies occur in f-block atoms, as well.

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• MAGNETISM
• magnetite--Fe3O4, natural magnetic oxide of iron
• 1600--William Gilbert concluded the earth is also
  a large spherical magnet with magnetic south at
  the north pole (Santa's habitat).
• NEVER FORGET opposites attract likes repel
• PARAMAGNETISM AND UNPAIRED ELECTRONS
• diamagnetic--not magnetic magnetism dies in
  fact they are slightly repelled. All electrons
  are PAIRED.
• paramagnetic--attracted to a magnetic field lose
  their magnetism when removed from the magnetic
  field HAS ONE OR MORE UNPAIRED ELECTRONS

• ELECTRON SPIN
• 1920--chemists realized that since electrons
  interact with a magnetic field, there must be one
  more concept to explain the behavior of electrons
  in atoms.
• ms--the 4th quantum number accounts for the
  reaction of electrons in a magnetic field

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• What about ferromagnetic?

• clusters of atoms have their unpaired electrons
  aligned within a cluster, clusters are more or
  less aligned and substance acts as a magnet.
  Don't drop it!!
• When all of the domains, represented by these
  arrows are aligned, it behaves as a magnet. This
  is what happens if you drop it! The domains go
  indifferent directions and it no longer operates
  as a magnet.

(Taken from summary notes posted on M. Brophys
website)
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Activities and Problem set __

• Ch 6 Problems Include pages from e-text write
  out answers show work
• First carefully study the sample exercises in
  chapter 6 (you dont have to copy them out) and
  then DO all in-chapter practice exercises
  according to the directions above..
• Do all GIST, and Visualizing concepts,
  problems
• end of chapter 6 exercises _________

• TB ch. 6 all sections required for SAT II and
  AP exams and most are required for regents exam
• View and take notes on the recommended animations
• POGIL activities on (1) Analysis of Spectral
  Lines and (2) Interaction of Radiation and Matter
  
• Online practice quiz due by ______