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Unit 6 : Quantum Mechanics, Molecular S tructure, and Orbital theory

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Unit 6: Quantum Mechanics, Molecular Structure, and Orbital theory

By: Eddie Yokana

and

Jake Gold

- Note double and triple bonds are counted as ONE electron domain.
- Electrons will always situate themselves to minimize repulsion.

- A bond is polar with electronegativity difference 0.5 or greater on the Pauling scale.
- All ionic bonds are polar.
- Fluorine and oxygen are polar with all atoms (besides itself).
- Generally if more than two spaces apart on periodical table then polar.

- A molecule is polar if polar bonds are asymmetrical.
- Naturally symmetrical molecules: linear, trigonal planar, tetrahedral, trigonalbipyramidal, octahedral, and square pyramidal.
- In a polar bond, the atom that originally had a higher electronegativity has a partial negative charge (δ-), and the atom with the lower electronegativity has a partial positive charge (δ+).

- There are six types of electromagnetic radiation: gamma rays, x-rays, visible light, infrared, and radio waves (in increasing wavelength).
- Visible light has wavelengths between 400nm to 700nm.
- Remember ROYGBIV- Red, Orange, Yellow, Green, blue, indigo, and violet (Order of light in decreasing wavelength.

- Light (c) goes at the speed of 2.9979 X 108 m/s through space, but slightly slower in air and about 1.5 times slower in water.
- C=fλ
- this equation is the relationship between frequency (f) and wavelength (λ).

- E=hf
- This equation solves for the energy (J) contained one photon. h is Plank’s constant, which is 6.626 X 10-34and f is frequency.

- KE= energy of a photon- energy threshold.
- The kinetic energy of an electron will equal the amount of energy hit by a photon minus the amount of energy it takes to emit the electron.

- KE= ½ mv2
- this equation can be used to solve for the velocity (v) of an electron if you know the kinetic energy (KE) of the electron. m is a constant, which is the mass of an electron, which is 9.11 X 10-34Kg.

- λ= h/(mv)
- is the de Broglie wavelength.

O2

- Bonding nodes- help form bond
- anti-bonding nodes- break apart bond
- Energy increases going up the chart
- Net Sigma bonds = (# of binding electrons in sigma bonds - # of anti-binding electrons in sigma bonds) / 2
- Net pi bonds = (# of binding electrons in pi bonds - # of anti-binding electrons in pi bonds) / 2
- Bond order = (# of bonding electrons - # of anti-bonding electrons) / 2
- Or = net sigma bonds + net pi bonds

Lewis Electron-Dot Example

STEPS

34e- - 8e- = 26e- remaining

6e-

7(4)e-

6+28

=34e-

SF4

6(4)e-= 24 e- needed

26e- > 24 e-

- Sum up valence electrons
- Make a basic single bond skeleton
- Determine electrons remaining (subtract those used in bonds)
- Determine electrons needed to complete octets
- Remain = needed : Finished
- Remain > needed : extra lone pair(s) on central atom
- Remain < needed : add extra pi bond for every 2 electron deficit

F

S

F

F

F

- If multiple isomers for molecule, then molecule with least formal charges will be the most stable.
- If two structures have the same amount of formal charges, then the more electronegative atom will form the bond.
- If both atoms are the same, then it could be a resonance molecule. This a way of describing delocalized electrons within a molecule

O

S

O

O

S

O

O

S

O

Periodic Trends

*Unequal electron affinity (i.e. >0.5) leads to polar bonds but the polar bonds must be asymmetrical for a molecule to be polar

Increasing Atomic Radius (size)

Increasing Electronegativity

Electron affinity is the ability of an atom to attract electrons from a bond.

*Since noble gases do not bond with other elements naturally, they are not included as we consider electronegativity.

Increasing Atomic Radius (size)

Increasing Electronegativity

- Note:
- As Electronegativity INCREASES, atomic radius DECREASES
- (Except Noble Gases)*Remember effective nuclear charge increases going across and orbitals are added going down

Electron Orbitals

“s” Block

“p” block

“d” Block

“f” Block

http://www.shs.d211.org/science/faculty/hlg/e%20conf%20travis/electron_configuration.htm

Electron Configuration

3 Ways of writing electron configurations:

Orbital box Notation

Spectroscopic Notation

Noble gas core Notation

*Note: The “d” orbitals and the “f” orbitals have different principle quantum numbers

http://www.shs.d211.org/science/faculty/hlg/e%20conf%20travis/electron_configuration.htm

Electron Configuration

Orbital box Notation:

3 Ways of writing electron configurations:

Orbital box Notation

Spectroscopic Notation

Shortcut Notation

- Spectroscopic Notation:

Noble gas core Notation:

- Four quantum numbers: n, l, ml, ms.
- N is the principle quantum number, and it identifies the electron shell or energy level (begins at 1).
- l describes the subshell (s,p,d,fusing the numbers 0,1,2,3 respectively)
- mldescribes the specific orbital within the subshell (For d’s subshell it goes -2,-1,0,1,2)
- msdescribes the spin. The spin is either +1/2 or -1/2. If electron is the first to go in a specific orbital, msis positive, and the second electron is negative.

WORKS CITED

- “Electron Density and Molecular Geometry”http://employees.csbsju.edu/hjakubowski/classes/ch123/Bonding/vsepr.gif
- “Low Energy to High Energy=Order of Filling”
http://www.chemistryland.com/CHM130W/10-ModernAtom/Spectra/ModernAtom.html

http://www.concord.org/~ddamelin/chemsite/d_bonding/configs.html

http://www.shs.d211.org/science/faculty/hlg/e%20conf%20travis/electron_configuration.htm