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MoL lecture summary and ppt summary
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MoL lecture summary and ppt summaryLecture 1
Nucleus (protons + neutrons) & electron cloud
Protons = positive Neutrons = neutral electrons = negative
Atoms are typically neutral #protons = # electrons
Electrons make up most of the volume -> important for molecule formation
Periods are horizontal & groups are vertical
Bohr’s model -> electrons in shells around the core
1st – 2 e
2nd – 8 e
3rd – 18 e
4th – 32 e
Electrons have wave-like properties
Shells contain subshells = atomic orbitals (AOs)
First shell s – number of atomic orbitals 1 – max of 2 electrons
Second shell s, p – number of atomic orbitals 1, 3 – max electrons 2 + 6 = 8
Third shell s, p, d – number of atomic orbitals 1, 3, 5 – max electrons 2+6+10=18
Core electrons – in inner shell (s)
Valance electrons – in outmost shell -> important for molecule formation
Relative energy of atomic orbitals
1s < 2s < 2p < 3s < 3p < 3d
Ground-state electron configuration – 3 rules
1. Aufbau principle: an electron goes into the atomic orbital with the lowest energy.
2. Pauli exclusion principle: no more than two electrons can be in an atomic orbital.
3. Hund’s rule: an electron goes into an empty degenerate orbital rather than pairing up.
Label axis! It gets you points in the exam
,Octet rule – atom is most stable if the outer shell is filled or empty!
Noble gasses – full shell
Ionic interactions
Sharing electrons results in covalent bonds
Polar and apolar covalent bonds
Apolar, delta – and delta +, where a part pulls more on the shared electrons
Electronegativity scale:
Greater dipole moment = more polar
Lone pairs are shown with dots
Lewis structures
To see how many electron pairs there need to be: see how many electrons it has in the outer shell
and see how many it needs for octate rule. Subtract these from each other this gives the amount of
bonds.
Skeletal representations
Atomic orbitals have geometrical shapes -> probability functions
S orbitals are spheric
,Two s orbitals form a sigma bond:
Orbitals are conserved – molecular orbitals (MOs)
# of atom orbitals combined = # of molecular orbitals
Atomic orbitals have geometrical shapes
P orbitals resemble dumbbells
Two p orbitals form a pi bond
Pi – bonding molecular orbital at the bottom & pi* antibonding molecular orbital at the top.
Nodal plane
Valance-shell electron pair repulsion – VSEPR
Bonding electrons and lone-pair electrons around an atom are positioned as far apart as possible!
Neither MOs nor VSEPR explains why all bonds are identical in methane (CH4)?
Hybrid orbitals result from combining AOs
, Promotion: 1e- of 2s orbital is promoted to the empty 2p orbital
Hybridization: s and p AOs are combined (=4 x sp^3 orbitals)
VSEPR + hybr. Orbitals explain the structure of CH4
Methane is tetrahedral. The tetrahedral bond angle is 109.5%
Lecture 2
The s orbitals from 4 H atoms each overlap with one sp3 orbital to form 4 sigma bonds
Sp3 – Sp3 -> C-C sigma bond
c-c pi bond originates from a p-p overlap
<- in purple is pi bond -> pi bond is formed by overlap of p&p
A pi bond cannot rotate!!!
Hybridization: 1s and 1 p AOs are combined (= 2x sp orbitals)
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