 DOR: Average Atomic Mass 9/17 (4 th /5 th ) 1)A gaseous element has two isotopes: G-102 with an atomic weight of and G-108 with an atomic weight of The percent abundance of the heavier isotope is %. What is the average atomic weight of element G ?

(Chapter 4)

 Light  Composed of small energy packets (photons)  Quantum = minimum amount of energy lost/gained by atom  Atoms can absorb or give off this energy

 Energy States in an Atom  Atoms can gain or loss energy.  Specific energy states within an atom.  Can be counted  Ground State = lowest energy state  Excited State = higher energy level than ground, gained energy

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 So, where does the Bohr Model fit in?  Electrons orbit around the nucleus at different energy levels/orbits.  Electron’s energy level = orbit level where electron is located.  Light absorption = electron moves from a state of low energy to high energy. “becomes excited”  Light Emitted = electron falls from an “excited” state of energy to a lower energy level.

 Main Energy Levels/Electron Shell  n = 1  Holds 2 electrons  n = 2  Holds 8 electrons  n = 3  Holds 18 electrons  n = 4  Holds 32 electrons

 Energy sublevels  Within the main energy level.  S = 1 orbital, can hold 2 electrons  p = 3 orbitals, can hold 6 electrons  d = 5 orbitals, can hold 10 electrons  f = 7orbitals, can hold 14 eletrons

 Back to the Bohr Model !

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 Example 1: He

 Example 2: F

 Ex. 3: Li

 Classwork: Bohr Models  PBe  NeS  OC  Na Mg **Read over “Hog Hilton” lab activity

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 1)Draw the Bohr Model for Carbon 2)The ________ state is the lowest energy state for an electron. 3)One electron of an atom is found at n=1 and another electron of the same atom is located at the n=3 energy level. Which electron has the highest energy? DOR: Bohr Model/Energy Levels

  Treats electron’s location as wave property  Defined by quantum numbers  Orbitals have different energies  Quantum numbers  Provide information about size, shape, and orientation of atomic orbitals  Define atomic orbitals from general to specific Quantum Theory

 Quantum Mechanical Model  Opposite charges attract, electrons are attracted to the nucleus of an atom  Takes a LOT of energy to keep electrons away from the nucleus.  Electrons are found at differing lengths from the nucleus and can only be present in certain locations

  Determines orbital size and electron energy  Same as “n” value/orbital in Bohr model  Positive whole number, NOT 0  Shells – orbitals with same value  n = 1, 2, 3, 4, etc. Principal Quantum Number (n)

  Defines orbital shape for a particular region of atom  Think of as “subshell”  Energy sublevels—within the main energy level  s = 1 orbital, can hold 2 electrons  p = 3 orbitals, can hold 6 electrons  d = 5 orbitals, can hold 10 electrons  f = 7 orbitals, can hold 14 electrons  Orbital Angular Momentum Quantum Number (l)

 Energy levels and Sublevels

  2p  4f How do you specify orbitals?

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  Describes the motion of an electron, spinning  As electron moves, magnetic field induced  Electrons with opposite spins, cancel magnetic field of other  Values: +1/2, -1/2 Electron Spin

 PP What does atomic structure REALLY look like?

  Quantum Worksheet Homework

 How are electrons distributed in an atom?

  Shorthand method for representing electrons’ distribution in orbitals within subshells  All orbitals have the same energy level— digenerate  Orbitals – mathematical expressions of probability of electron’s location  Electrons occupy orbitals in a way that gives LOWEST energy state Electron Configuration

  Visual representation of electron configuration  Represents electrons’ spins ( ,  ) Orbital Diagrams

  Electrons occupy the LOWEST energy orbital available  Lazy Hogs ! Aufbau Principle

  Developed by Friedrich Hund  Creates the most stable electron arrangement  Based on electron spin Hund’s Rules

 1)One electron MUST occupy each orbital BEFORE electrons are paired in the same orbital. 2)Electrons added to subshell with the same spin (+1/2, -1/2) so each orbital has one electron. Hund’s Rules cont.

  Only 2 electrons occupy each orbital  Electron spins MUST be opposite/paired when 2 electrons occupy the same orbital  +1/2, -1/2  Pauli Exclusion Principle

  Period numbers = principal quantum number of valence shell electrons  Subshells fill with electrons at different regions within periodic table (s section, p section) Using the periodic table- -

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 Ex. 1 Nitrogen

 Ex. 2 Cr

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 1)Ca 4) O 2)P5) Li 1)Mn Homework Practice