Chemistry I Quantum Mechanics Notes by ixp95504

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									Chemistry I: Quantum Mechanics Notes

Bohr Model of Atom:
    • electrons move around nucleus in orbits similar to how planets orbit the sun
    • energy levels for electrons are quantized

Major developments that put Bohr’s Model into question:

Einstein: Light energy exhibits properties of matter. Matter and energy are different forms of the same thing.

De Broglie: Electrons move around nucleus in waves.

Heisenberg: Uncertainty Principle: it is impossible to measure the momentum (velocity) and location of an electron at
the same time.

Max Born: Interpreted an equation discovered by Schrodinger to mean that although we cannot predict the exact
location and path of an electron, we can describe its path in terms of probability. This is the fundamental concept
behind Quantum Mechanics.

         Bohr Model of Atom:                                     Quantum Mechanical Model of Atom:




How is Quantum mechanics similar to Bohr Model?

         1. There are quantized energy levels.



         2. Electrons jump from level to level when excited.



How is Quantum Mechanics different from Bohr Model?

         1. We can describe an electron’s path in terms of probability, not precise orbits.



         2. Each energy level (shell) is divided into sublevels (subshells).



         3. Each sublevel is made up of 1 or more orbital.



         4. Each subshell corresponds to a different shape of the orbital. (Not always spherical!)
Electron Configuration Notes

Important: As you increase in atomic number, each atom adds a new electron:

       •   Every successive element has the same electron configuration as the previous element, plus the additional
           electron. (Some exceptions)

       •   The energy level (shell) of this electron is described by the row where the atom is on the periodic table.
           (Exceptions (which won’t make sense to you now, but they will someday): “d” subshell is always 1 level
           behind; “f” subshell is always 2 levels behind.)

       •   The sublevel (subshell) of this electron is described by the group of columns where the atom is on the periodic
           table.

Energy Levels are represented by numbers: 1, 2, 3, 4…

Subshells are represented by letters: s, p, d, f

Examples:

 H:                                                                Si:


 He:                                                               P:


 Li:                                                               S:


 Be:                                                               Cl:


 B:                                                                Ar:


 C:                                                                K:


 N:                                                                Ca:


 O:                                                                Sc:


 F:                                                                Ti:


 Ne:                                                               Zn:


 Na:                                                               Ga:


 Mg:                                                               Kr:


 Al:
Orbital Notes

Remember how each energy level is divided into subshells?

Subshells are also divided!

Subshells are divided into orbitals.

Each orbital contains 2 electrons. (of opposite spins)

So, in summary:

____________________are divided into ____________________, which are divided into ____________________,
each of which contains ____________________.

Each type of subshell contains a certain number of orbitals:

             subshell         # of orbitals              # of electrons
             s
             p
             d
             f


Hund’s Rule: When filling a subshell, electrons enter each orbital until all orbitals contain one electron with their
spins parallel.


ORBITAL BOX DIAGRAM.
This uses a little circle, O, to represent each orbital. We group the orbitals of a subshell together, and label them. An
orbital with no electrons in it (an “empty house”) is left as a circle: O. With one electron we put a diagonal line in to
show one electron in the orbital.
The line can slant either way; one way, it's spinning one direction, the other slant represents an electron spinning the
other way; W or Ø . Two electrons in an orbital need both diagonal lines. U Let's take an example; Na, sodium,
element #11.

Electron Configuration:

1s 2 2s 2 2p 6 3s 1

Its orbital box diagram looks like this:

1s           2s         2p              3s
U            U          UUU             Ø


Do other examples!
Probability Diagram Notes

Most of the orbitals for a particular subshell are shaped the same:

         subshell           shape



         s



         p



         d



         f




Draw probability diagrams for Na and Mg:




Bonding Electrons (also called valence electrons): Electrons involved in bonding

The electrons in the s and p subshells of the highest energy level.

Examples:
How many bonding electrons are there Na? Mg?




Lewis Dot diagrams : Represent the bonding electrons
Stability

Rule:

     •      atoms/ions are stable when ½ of a subshell is filled
     •      atoms/ions are even more stable when a subshell is entirely filled
     •      atoms/ions are the most stable when the valence shell (valence electrons) are entirely filled (Octet Rule)

Examples of atoms with unusual electron configurations:

Cr


Zn



More Examples:
Use electron configuration and orbital box diagrams to explain why:

     •      Noble gases like Kr do not form ions




     •      Nitride is a 3- ion




     •      Sulfide is a 2- ion




     •      Iodide is a 1- ion




     •      Potassium is a 1+ ion




     •      Magnesium is a 2+ ion




     •      Aluminum is a 3+ ion

								
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