Chapter 10 Part One Modern Atomic Theory

1
Chapter 10- Part OneModern Atomic Theory

• Objectives
• Review history (10.1)
• Describe electromagnetic radiation (10.2)
• Describe the Bohr atom (10.3)
• Explain energy levels of electrons and diagram
  atomic structures for elements (10.4 10.5)

2
Review

• Dalton
• Thomson
• Rutherford
• Model doesnt explain how the negative electron
  can stay in orbit and not be attracted to the
  positive proton

3
Electromagnetic Radiation

• Light travels in waves, similar to waves caused
  by a moving boat or a pebble tossed in a pond
• Light is a form of Electromagnetic Radiation
• Form of energy that exhibits wavelike behavior as
  it travels through space

4
Electromagnetic Radiation

• All waves can be described in 4 ways
• Amplitude the height of the wave, results in
  the brightness or intensity of the light
• Wavelength (l) distance between consecutive
  peaks in a wave
• Frequency (n) number of waves that pass a
  given point in a second

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Electromagnetic Radiation

• Speed of light in air Electromagnetic radiation
  moves through a vacuum at speed of 3.00 x 108 m/s
• Since light moves at constant speed there is a
  relationship between wavelength and frequency
• c ln
• Wavelength and frequency are inversely
  proportional

6
Electromagnetic Spectrum
7
Quantum Theory

• Wave theory does not explain
• Heated iron gives off heat
• 1st red glow yellow glow white glow
• How elements such as barium and strontium give
  rise to green and red colors when heated

8
Quantum Theory

• Max Planck (1858-1947)
• Proposed that there is a fundamental restriction
  on the amounts of energy that an object emits or
  absorbs, and he called each of these pieces of
  energy a quantum.
• Energy is release in Quanta

9
Quantum Theory

• A quantum is a finite quantity of energy that can
  be gained or lost by an atom
• E hn E energy
• v frequency
• h 6.626 x 10-34 J/s
• This constant, h, is the same for all
  electromagnetic radiation

10
Photoelectric Effect

• The emission of electrons by certain metals when
  light shines on them
• Albert Einstein (1905) used Plancks equation to
  explain this phenomenon
• proposed that light consists of quanta of energy
  that behave like tiny particles of light
• Photon individual quantum light (also known as
  a particle of radiation)

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Photoelectric Effect

• He (Einstein) explained that the photoelectric
  effect would not occur if the frequency and
  therefore the energy of each photon is too low to
  dislodge an electron.
• Analogy
• 70 cents placed in soda machine no soda
• 30 cents more and you will get your soda

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Now

• Light can be described as both particles and
  waves
• Dual Wave-Particle Nature of Light was accepted
• What does this mean for the atom???

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LineSpectrum

• Elements in gaseous states
  give off colored light
• High temperature or high voltage
• Always the same
• Each element is unique
• http//home.achilles.net/jtalbot/data/elements/

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Line Spectrum

• Ground state
• Lowest energy level available
• Excited state
• State in which electron has a higher potential
  energy than in its ground state
• Farther from nucleus
• Higher potential energy

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Line Spectrum

• Electron falls from higher energy level to lower
  oneemits light at a specific frequency
• Color of light emitted depends on difference
  between excited state and ground state
• See figure 10.5 page 201

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Line Spectrum

• Each band of color is produced by light of a
  different wavelength
• Each particular wavelength has a definite
  frequency and has definite energy
• Each line must therefore be produced by emission
  of photons with certain energies

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Line Spectrum
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Line Spectrum

• Whenever an excited electron drops from such a
  specific excited state to its ground state (or
  lower excited state) it emits a photon
• The energy of this photon is equal to the
  difference in energy between the initial state
  and the final state.

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Niels Henrik David Bohr

• 1885-1962
• Physicist
• Worked with Rutherford
• 1912
• Studying line spectra
• of hydrogen

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Niels Henrik David Bohr

• 1913 proposed new atomic structure
• Electrons exist in specific regions away from the
  nucleus
• Electrons revolve around nucleus like planets
  around the sun

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The Bohr Atom

• Nucleus with protons and neutrons
• Electrons move in stationary states which are
  stable (paths or orbits)
• When an electron moves from one state to another
  the energy lost or gained is done is ONLY very
  specific amounts
• Each line in a spectrum is produced when an
  electron moves from one stationary state to
  another

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The Bohr Atom

• Model didnt seem to work with atoms with more
  than one electron
• Did not explain chemical behavior of the atoms

23
Wave Matters

• Louis de Broglie (1924)
• Proposed that electrons might have a
  wave-particle nature
• Used observations of normal wave activity

24
Wave Matters

• Erwin Schrodinger (1926)
• Used mathematical understanding of wave behavior
  devised an equation that treated electrons
  moving around nuclei as waves
• Quantum Theory

25
Quantum Theory

• Describes mathematically the wave properties of
  electrons and other very small particles
• Applies to all elements (not just H)

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Energy Levels of Electrons

• Principal energy levels
• Designated by letter n
• Each level divided into sublevels
• 1st energy level has 1 sublevel
• 2nd energy level has 2 sublevels
• Etc.

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Energy Levels of Electrons
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Orbitals

• Electrons dont actually orbit like planets
• Orbital region in space where there
  is a high probability of finding a given electron
• Each orbital sublevel can hold 2 electrons

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Orbitals
Each sublevel (orbital) has a specific shape
http//daugerresearch.com/orbitals/
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Orbitals

• Pauli exclusion principle an atomic orbital can
  hold a maximum of two electrons which must have
  opposite spins
• Electrons can only spin in two directions
• Shown with arrows

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Rules for Orbital Filling

• Paulis Exclusion Rule
• No two electrons have the same set of quantum
  numbers
• Hunds Rule
• Electrons will remain unpaired in a given orbital
  until all orbitals of the same sublevel have at
  least one electron
• 1s 2s 2p 3s 3p

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Rules for Orbital Filling

• Diagonal Rule
• The order of filling once the d f
  sublevels are being filled
• Due to energy levels

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Rules for Orbital Filling
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Quantum Numbers

• Numbers that specify the properties of atomic
  orbitals and their electrons
• Principle Quantum Numbers
• Symbolized by n, indicates the main energy levels
  surrounding a nucleus, which indicates the
  distance from the nucleus (shells or levels)

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Quantum Numbers

• Orbital Quantum Number
• Indicates the shape of an orbital
• (subshell or sublevels)
• s, p, d, f
• Principal Quantum Orbital Quantum
• 1 1s
• 2 2s, 2p
• 3 3s, 3p, 3d
• 4 4s, 4p, 4d, 4f

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Quantum Numbers

• Magnetic Quantum Number
• Indicates the orientation of an orbital about the
  nucleus
• Orbital position with respect to the
  3-dimensional x, y, and z axes

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Quantum Numbers

• Spin Quantum Number
• Indicates two possible states of an electron in
  an orbital
• Type of Orbital Number of Orbitals
• s 1 ( )
• p 3 (x, y, z) ( , , ,)
• d 5 ( , , , , )
• f 7
• Each orbital holds a maximum of 2 electrons

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Application of Quantum Numbers

• Several ways of writing the address or location
  of an electron
• Lowest energy levels are filled first
• Electron Configuration using the diagonal rule,
  the principal quantum number (n), and the
  sublevel write out the location of all electrons
• 12C
• 32S

1s22s22p2
1s22s22p63s23p4
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Application of Quantum Numbers

• Orbital filling electron diagram using Hunds
  rule and the diagonal rule write out the location
  of all electrons
• See examples on whiteboard

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Homework

• Worksheet 1
• Question 11
• Paired Exercises 27-33 odd
• Additional Exercises 54

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Chapter 10 Part TwoThe Periodic Table

• Objectives
• Understand the arrangement of the Periodic Table
  (10.6)
• Identify connections between electron
  configuration and placement on the periodic table

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The Periodic Table

• 1869 arrangement proposed by Dmitri Mendeleev
• And Lothar Meyer (different layout)
• Still similar today
• Based on increasing atomic masses and other
  characteristics
• Was able to predict properties of elements not
  yet discovered.and was correct!

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The Periodic Table

• Horizontal rows
• Periods
• Corresponds to outermost energy level
• Vertical Columns
• Groups or families
• Similar properties reactions

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The Periodic Table

• Several systems for naming groups
• Left to right, 1-18
• Roman numerals and A and B
• Used in this book
• Group A Representative Elements
• Noble Gases
• IA Alkali Metals
• IIA Alkaline Earth Metals
• VIIA - Halogens
• Group B Transition Elements

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The Periodic Table

• Chemical behavior and properties of elements in a
  particular family similar
• Have the same outer shell electron configuration
• Figure 10.15 page 211
• Noble gas configuartion (shortcut)
• Use previous noble gas in square brackets
• Finish with valence electrons

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The Periodic Table

• Examples
• K is 1s22s22p63s23p64s1 or Ar4s1
• Ca is 1s22s22p63s23p64s2 or Ar4s2
• Write abbreviated configuration for the following
  elements
• Fr
• Y

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The Periodic Table

• Arrangement of Periodic Table also means that
  elements filling similar orbitals are grouped
• s block
• p block
• d block
• f block
• Know these blocks

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The Periodic Table - Highlights

• The number of the period corresponds to the
  highest energy level occupied by electrons in
  that period
• The group numbers for the representative elements
  are equal to the total number of valence
  electrons in that group

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The Periodic Table - Highlights

• The elements of a family have the same outermost
  electron configuration
• (just different energy levels)
• The elements within each of the s, p, d, and f
  blocks are filling the corresponding orbitals
• There are some discrepancies with order of
  filling
• (not covered in this book)

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Homework

• Worksheet 2
• Questions 12-16 even
• Paired Exercises 39-44 all 49-52 all
• Additional Exercises 57 59
• STUDY FOR TEST ON THURSDAY
• Questions answered from 500 -545 pm
• Test starts at 545 pm
• Chapter 11 Lecture will begin at 700 pm