Atomic Theory

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Unit 2

• Atomic Theory

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Video links

• overview of atomic history
• http//www.youtube.com/watch?featureplayer_detail
  pagevk1RHY8QcN1s

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I. Atomic History

• A. The Greeks
• Democritus Philosopher
• All matter is made of tiny, indivisible parts
  called atoms
• Developed word atomos meaning not divisible

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John Dalton (1803-1808)

• Used experiments with gases to develop the
  Atomic Theory
• Determined atoms looked like cannonballs or
  solid masses

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Daltons Atomic Theory

• 1) All elements are made of atoms
• 2) Atoms of each element are all the same, or
  have the same masses
• 3) Atoms of different elements are different, or
  have different masses
• 4) Atoms cannot be created or destroyed
• 5) Atoms combine in small, whole number ratios

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J.J. Thomson (1897)

• Developed Cathode Ray experiment
• Said atoms consisted of particles smaller than an
  entire atom
• Discovered that the smaller particles within an
  atom had a negative charge
• Discovered 1st subatomic particle Electron
• Founded Plum Pudding Model Electrons were
  embedded within a positively charged mass

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Cathode Ray Tube Experiment
Thomson manipulated cathode rays with a magnet to
discover that subatomic particles existed and
that they had negative charges
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Ernest Rutherford (1898)

• Discovered alpha and beta radiation emitted from
  certain radioactive substances
• Developed and used Gold Foil Experiment
• First to separate the smaller parts of the atom

• Discovered the nucleus
• Placed electrons outside the nucleus
• Stated that atoms are composed of lots of empty
  space

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Rutherfords Gold Foil Experiment
micro.magnet.fsu.edu/electromag/java/rutherford/

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Niels Bohr (1922)

• Bohr analyzed work of others and studied atomic
  spectra, or light, given off by the elements
• Described the Atomic Spectra of elements
• Developed Solar System model
• Moved electrons from single, giant pathway into
  discrete energy levels around the nucleus
• Each energy level contained 2, 8, 18, 32, etc.
  electrons total

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Bohr Model of the Atom

• Stated that electrons moved around the nucleus in
  orbits or energy levels
• As electrons gain energy, they jump up energy
  levels, then release this energy to generate
  spectra

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Bohrs model and the atomic spectrumhttp//jerse
y.uoregon.edu/vlab/elements/Elements.html

• The spectral lines in the visible region of the
  atomic emission spectrum of barium are shown
  below.
•                                                 
                                                    
                                                    
                                                    
                    
• Spectral lines exist in series in the different
  regions (infra-red, visible and ultra-violet) of
  the spectrum of electromagnetic radiation.
• The spectral lines in a series get closer
  together with increasing frequency.
• Each element has its own unique atomic emission
  spectrum.

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Erwin Schrodinger (1930)

• Developed mathematical equations representing
  electrons
• Electrons had wave and particle behaviors
• Created Wave-Mechanical or Modern model
• Most scientists use this model today
• Placed electrons in orbitals

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Electron Cloud Model

• Created paths for electrons within Bohrs energy
  levels
• Only 2 electrons per path
• Electron paths, or ORBITALS, are mathematical
  equations describing probability densities for
  electrons
• Developed sublevels with discrete paths within
  each energy level

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II. Subatomic Particles

• Particles
• 1) Protons
• found in the nucleus of an atom
• charge of 1, mass of 1.0073a.m.u.
• 2) Neutrons
• found in the nucleus of an atom
• no charge, mass of 1.0087a.m.u.

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A. Subatomic Particles

• 3) Electrons
• Found outside the nucleus in regions of
  probability orbitals
• Charge of 1, mass of 5.46 x 10-4 a.m.u., or
  1/1836 a.m.u.
• Have particle and wave properties

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B. Atomic Number

• Atomic number the number of protons in the
  nucleus
• All atoms of the same element have the same
  atomic number

• Atoms arranged on PT by increasing atomic numbers
  
• In neutral atoms
• Atomic number equals number of electrons

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C. Isotopes

• Isotopes atoms of the same element that have
  differing numbers of neutrons in their nucleus,
  different mass number, but same atomic number
• Same number of protons!!!
• Changing number of neutrons affects properties
  radioactivity

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D. Atomic Mass

• Atomic Mass Number number of protons plus the
  number of neutrons in the nucleus
• Whole number!!

• Mass Number changes when using different isotopes
  
• Written in isotopic notations, just subtract the
  top from bottom values

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E. Ions

• Ions atoms of the same element that have lost
  or gained electrons
• Have overall () or (-) charge
• Same numbers of protons, number of neutrons
  irrelevant
• Positive ions have LOST electrons
• Negative ions have GAINED electrons

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F. Atomic Mass (average)

• Atomic Mass weighted average of the natural
  isotopes times their percent abundance
• Decimal value on PT
• Accounts for the natural existence of various
  isotopes
• Ex calculate the atomic mass of carbon given
  that 98.92 is carbon-12 and 1.108 is carbon-13

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Virtual textbook

• http//www.chem1.com/acad/webtext/intro/int-1.html
  SEC1

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III. Electronic Structure
A. EMS Electromagnetic Spectrum
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A. EMS Electromagnetic Spectrum

• EMS continuous series of various types of
  energy, separated by their wavelengths and
  frequencies
• Visible light small portion only part we can
  see without instruments
• Continuous spectrum picture of all colors of
  visible light as they pass through a prism

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EMS continued

• Wavelength distance between 2 peaks or troughs
  of 2 consecutive waves
• Symbol ?
• Greek letter lambda
• Units are usually in m or nm

• Frequency the number of peaks or troughs that
  pass a single point in one second
• Symbol ?
• Greek letter nu
• Units are usually in 1/s or s-1 or Hz

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Calculations using lambda and nu

• c ??
• C speed of light
• C 3.0 x 108 m/s
• E h?
• E energy of photon
• h Plancks constant
• h 6.63 x 10-34 Js

• All electromagnetic radiation travels at the
  speed of light
• Can calculate the energy of the
  radiation/electron given the wavelength

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Plancks Constant

• Planck observed hot, glowing matter
• Concluded different substances glow different
  colors at different temperatures
• Determined matter releases energy in tiny,
  discrete packets called quanta
• Developed constant to relate energy and
  temperature, Plancks constant, h
• h 6.63 x 10-34 Js

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Light traveling as waves
All colors of light energy travel at the same
speed, just different wavelengths!
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Particle vs. Wave Behavior of Light
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Wave behavior of light
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B. Photoelectric Effect

• Einstein used Plancks idea of quanta and photons
  to describe the photoelectric effect
• Light of a certain wavelength shines on clean
  metal, causing the metal to eject electrons

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Bohrs Model conclusions made

• Bohr used the idea of quanta to explain the
  bright-line emission spectra
• Stated that each elements atomic spectrum is
  unique
• Electrons exist in ground state energy levels, as
  listed via the periodic table

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Bohr Model of the Atom

• Stated that electrons moved around the nucleus in
  energy levels
• Electrons will gain and lose energy at will
• This generated the elements atomic spectrum

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Bohrs model
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Useful Websites and References

• //www.avogadro.co.uk/light/bohr/spectra.htm
• shows formation of spectral lines for hydrogen
• idea of ground vs. excited state
• //jersey.uoregon.edu/vlab/elements/Elements.html
• Periodic table showing the absorption and
  emission spectra for each element
• Also check out Wikipedia under Bohr atom and
  Atomic spectra!

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Creation of an emission spectrum

• If electrons absorb packets of energy, quanta,
  they temporarily move to into a higher energy
  level, called the excited state
• The electrons then release this quanta of energy
  and fall back down to ground state
• The release of energy generates the bright-line
  emission spectrum

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Examples of Bohr Diagrams
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IV. Electron Configurations

• Energy Levels
• These are areas with a high possibility of
  finding electrons with similar potential energies
• 7 energy levels total

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Bohr Diagrams and Energy Levels

• Bohr Diagrams show the numbers of protons and
  neutrons in the nucleus
• Shows electrons in their respective energy levels
  

• Energy levels hold
• 1st holds 2 electrons
• 2nd holds 8 electrons
• 3rd holds 18 electrons
• 4th holds 32 electrons
• Etc..

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B. Sublevels

• Sublevels are divisions within each energy level
• Represent the shapes and orientation in 3D space
• Too many electrons within the energy levels
  they lose momentum and will crash into the
  nucleus--- not good!

• 1st energy level has 1 sublevel s
• 2nd has 2 sublevels s and p
• 3rd has 3 sublevels s, p, and d
• 4th has 4 sublevels s, p, d, and f

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Sublevels and Shapes

• s is spherical and has a max of 2 electrons
• p is dumbbell shaped and has a max of 6
  electrons
• d is cloverleaf shaped and holds up to 10
  electrons
• f is a split cloverleaf with a max of 14
  electrons
• http//micro.magnet.fsu.edu/electromag/java/atomic
  orbitals/index.html

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Order of Sublevel Filling

• It does not go in order
• 1s2
• 2s2 2p6
• 3s2 3p6 3d10
• 4s2 4p6 4d10 4f14
• 5s2 5p6 5d10 5f14
• 6S2 6P6 6d10
• 7s2 7p6

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Orbitals within Sublevels

• Each sublevel consists of 1 to 7 orbitals areas
  of probability for finding an electron
• Each path or orbital only holds 2 electrons

• The 2 electrons within in each orbital each have
  a different spin
• This allows the electrons to exist in the same
  area without conflicting

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Extended and Abbreviated Configurations

• Electron Configurations way to describe how the
  electrons are distributed around an atom and
  within the energy levels and sublevels
• Ground state configurations are same order as
  electrons on PT
• Excited state configurations have one electron
  shifted to a higher energy level

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Writing Electron Configurations

• Electrons add in the same order as the atomic
  numbers of the PT
• Aufbau Principle adding electrons in the exact
  order of the PT

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Writing Configurations

• Add in order of arrows for Neutral, Ground State
  atoms
• Examples

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Abbreviated Configurations

• Abbreviated configurations show only the
  placement of electrons added after the last
  noble gas
• Ex

• Bracket the configuration of the last noble gas
  group 18 and add remaining electrons
• Ex

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Orbital Notations and Rules

• Orbital notations are specialized versions of a
  full electron configuration showing the spin of
  each electron within an orbital
• Draw the orbitals present for each sublevel and
  fill with spin-paired electrons

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Rules for Configurations

• 1. Hunds Rule electrons in the p, d, and f
  sublevels must be added to each orbital first,
  before one flips to spin-pair and fill the
  orbital

• 2. Pauli Exclusion Principle no 2 electrons
  may be in the same orbital and have the same
  spin no 2 electrons will have the same 4 quantum
  numbers

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Rules cont

• Heisenbergs Uncertainty Principle states that
  the electrons momentum and position cannot be
  accurately determined at the same time
• Example

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Excited State vs. Ground State

• Excited State configurations show one electron
  has moved into a higher energy level, leaving an
  unfilled space below

• Ground state configurations are written in order
  of the periodic table
• Total of electrons Atomic
• for BOTH !!

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E. Lewis Dot Structures

• Lewis Dot Structures are pictures showing the
  placement and number of valence electrons for an
  element
• Structure
• s1s2
• p6 p1
• p3 p4
• p5p2

• Valence electrons are s and p outer shell
  electrons
• Maximum of 8!
• Ex

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F. Quantum Numbers

• Each electron in an atom is assigned a set of 4
  quantum numbers
• These numbers tell the exact address of an
  electron, regardless of the element
• No 2 electrons have the same 4 quantum numbers!

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Quantum Numbers

• 1 Principle Quantum Number (n)
• First number
• Represents the energy level of the electron
• Values range from
• 1 to 7

• 2 Azimuthal Spin Number (l)
• Second Number
• Represents the sublevel
• Describes the shape of the orbital
• Values from 0 to 3

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Quantum Numbers

• 3 Magnetic Spin Number (ml )
• Tells the orientation of the orbital along x, y,
  z axes
• Values for
• l 0, ml 0
• l 1, ml 1, 0, -1
• l 2, ml 2, 1, 0, -1, -2
• l 3, ml 3, 2, 1, 0, -1, -2, -3

• 4 Spin Number (ms)
• Tells if the electrons spin clockwise, or
  counterclockwise
• Values
• 1/2 spin up or
• 1/2 spin down