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

1. Unit 6: Quantum Mechanics, Molecular Structure, and Orbital theory By: Eddie Yokana and Jake Gold

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

3. Polarity • 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+).

4. Electromagnetic Radiation • 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.

5. Equations • 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.

6. Molecule Orbital Energy Levels 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

7. 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

8. Lewis Dot Structure continued • 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

9. 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

10. Electron Orbitals “s” Block “p” block “d” Block “f” Block

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

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

13. Quantum Numbers • 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.

14. 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