1.  DOR: Average Atomic Mass 9/17 (4 th /5 th ) 1)A gaseous element has two isotopes: G-102 with an atomic weight of 102.11 and G-108 with an atomic weight of 108.08. The percent abundance of the heavier isotope is 34.550%. What is the average atomic weight of element G ?

2. (Chapter 4)

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

4.  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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8.  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.

9.  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

10.  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

11.  Back to the Bohr Model !

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15.  Example 1: He

16.  Example 2: F

17.  Ex. 3: Li

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

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20.  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

21.   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

22.  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

23.   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)

24.   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)

25.  Energy levels and Sublevels

26.   2p  4f How do you specify orbitals?

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29.   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

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

31.   Quantum Worksheet Homework

33.  How are electrons distributed in an atom?

34.   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

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

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

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

38.  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.

39.   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

40.   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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42.  Ex. 1 Nitrogen

43.  Ex. 2 Cr

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