Transcript Slide 1

Kevin Mospan and Elizabeth Schultz
Unit 6 Review
Atomic and Molecular Structure
Lewis structures- know how to draw the dot structures
Molecular shapes
oDifferent shape names/molecular geometry
oBond angles (See Molecular orbital theory –other side)
Formal charge
oFormal charge closer to 0 is favorable
Bond polarity
oDifference in electronegativity 0.4<x<1.7
2
Molecular polarity
oSymmetrical bonds cancel
Electromagnetic radiation
o. = c = 2.99792458 x 108 m/s
oElectromagnetic spectrum
Energy quantization
oPlanck
oE=h. (h≈6.6260755x10-34 J.s)
Photoelectric effect
oWhen light strikes the surface of a metal and e- are ejected
oE=h.c/
Bohr model
oRadius of the circular orbits increase as n increases
oAn atom with its e- in the lowest possible energy levels is said to be in its “ground state”
oWhen an e- occupies an orbit greater than the lowest possible energy level it is said to be in
an “excited state”
oΔE=-Rhc(1/nf2 - 1/ni2)
Rhc=1312 kJ/mol
Wave/particle duality
oTaken from idea that light, usually considered to exhibit wave properties, actually consists of
particles (photons)
oSimilarly, massless particle, e-, can exhibit both particle and wave properties
DeBroglie wavelength
Value of l
o=h/m.
Corresponding subshell
oh=Planck’s constant
label
om=mass (kg)
0
s
o=velocity (m/s)
Quantum #s
1
p
on=principal quantum #
2
d
integers from 1-
ol=angular momentum quantum #
3
f
characterizes subshell of orbital
starts at 0 and goes to n-1
om l =magnetic quantum #
0, 1, 2,…,  l
# of values of ml for a given subshell is 2 l +1
oms= electron spin magnetic quantum#
 ½
Orbitals
os, p, d, f (…)
obattleship/revised periodic table method
os has 1 orbital
op has 3 orbitals
od has 5 orbitals
of has 7 orbitals
Quantum numbers for electron configurations
oDiamagnetic means not magnetic (diametrically opposed)
oParamagnetic means exhibits magnetism caused by unpaired eoPauli exclusion principle: no two e- can have the same set of 4 quantum numbers
(only two e- per box)
Electron configurations of atoms and ions
oHund’s rule: the most stable arrangement of e- is that with the maximum number of
unpaired eoNoble gas notation
oHalf filled orbitals are stable
Atomic and chemical periodic trends
(arrows indicate increase for all)
oRadius ↓
←
oElectronegativity ↑
→
oIonization energy ↑ →
Valence bond theory
oHybridization
Bond angles
oSigma and pi bonds
oSigma bond necessary for a pi bond to occur
Molecular orbital theory
oBond order =
½( #e- in bonding MOs - #e- in antibonding MOs)
oAnti-bonding –principle of molecular
orbital theory that isn’t in valence bond
theory, antibonding orbital are higher in
energy than bonding orbitals