Quick Review of the physics behind the Quantum
Mechanical Model
•Light consists of electromagnetic waves (energy waves).
•Amplitude: height of wave
•Wavelength(λ): distance between crests
•Frequency(v): the # of wave cycles per unit time SI: Hz (s-1)
•The electromagnetic spectrum includes: radio waves,
radar, microwaves, infrared, VISIBLE LIGHT, uv, x-rays,
gamma rays, and cosmic rays.
•All of these waves travel at the speed of light (c) in a
vacuum. c = 3.0x108 m/s
•Frequency (ν) and wavelength (λ) are inversely related.
•
ν = λ AND ν = λ
• c = ν λ the speed of light equals the product of
frequency and wavelength... so we can calculate ν if given
λ, or λ if given ν.
•Again, c = 3.0x108 m/s and the units of ν are Hz or s-1
•Example: Calculate the wavelength of the yellow light
emitted by a sodium lamp if the frequency of the radiation
is 5.10x1014 Hz.
Emission Spectra
•Sunlight (white light) consists of light with a range of
wavelengths and frequencies.
•A prism is used to separate different wavelengths into a
spectrum of colors.
•The ν and λ are CHARACTERISTIC of each color.
•Every element emits light when excited. The atom absorbs
E, then loses E as they emit light.
•Passing the light emitted though a prism gives the atomic
emission spectrum of the element.
•White light = continuous spectrum
•Excited atoms’s emission spectra = only a few lines
• Each line corresponds to one frequency of light emitted by
the atom.
•The emission spectrum of each element is unique.
•This is what we mimicked with the flame test.
Physicist Max Planck found the
amount of E absorbed or emitted is
proportional to the frequency (v)
E=vh
E= energy in Joules (J)
h= Planck’s constant 6.62x10-34 Js
(J times sec)
Einstein explained the
photoelectric effect, where metals
eject electrons when light hits
them.
Example: Calculate the energy in
Joules of a photon with a
frequency of 5.00x1015 s-1.
Bonding Theories
Valence-Shell Electron-Pair Repulsion
(VSEPR): because electron pairs repel,
molecular shape adjusts so the valence electron
pairs are as far apart as possible.
Hybridization: involves the overlap of atomic
orbitals, and gives information about bonding
and shape.
Hybridized Orbitals
When different orbitals mix together they form the
same number of hybrid orbitals.
electron
Linear
geometry--> (1 domain)
Hybridization
s
Linear (2
domains)
Trigonal
Planar
Tetrahedral
sp
2
sp
3
sp
DRAW THE STRUCTURES OF:
CH4
C2H4
C2H4
SIGMA BOND: A SINGLE BOND
PI BOND: DOUBLE BOND
Polar Bonds and Molecules
Covalent bonds involve sharing electrons around
atoms, but the sharing is not always equal.
There is a tug-of-war between the two atoms and
the shared pair of electrons
When like atoms pull equally, the bond is nonpolar
covalent.
Examples: O2, H2, Cl2, N2
Polar Bonds and Molecules
Covalent bonds involve sharing electrons around
atoms, but the sharing is not always equal.
There is a tug-of-war between the two atoms and
the shared pair of electrons
When like atoms pull equally, the bond is nonpolar
covalent.
Examples: O2, H2, Cl2, N2
When a covalent bond joins two different atoms and the
sharing is unequal, the bond is a polar covalent bond (or
just polar bond).
The unequal sharing is due to differences in
electronegativity.
The presence of a polar bond in a molecule usually
makes the entire molecule polar.
Examples: HCl (a polar bond), H2O (a polar molecule)
Polar molecules have a slightly positive and and a
slightly negative end. A molecules that has two poles
like is is called a dipole.
Is carbon dioxide a
polar molecule?
But what about water?
Using electronegativity
to determine bond type
Bond Type -->
Nonpolar
Covalent
Polar Covalent
Ionic
Difference in
Electronegativity
0-0.45
0.45-2.0
>2.0
Example: predict the bond type that will form
between each of the following pairs.
N (3.0) and H (2.1)
F (4.0) with itself
Ca (1.0) and Cl (3.0)
Intra- vs. Inter- molecular
forces
Intramolecular forces: forces holding molecules
together. Ex: ionic and covalent bonds
Intermolecular forces: attractions between
molecules.
Weaker than intramolecular forces
Determine solid/liquid/gas
Intermolecular Forces
Hydrogen Bonds
London Dispersion Forces
Dipole-Dipole Interactions
Ion-Dipole Interactions
Covalent Network Solids