Orbitals Flashcards

1
Q

promotion energy

A

energy required to promote e- from ground state to valence state (= one type of excited state configuration used for bonding)

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2
Q

localised

A

e- are confined to a particular bond or atom

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3
Q

Lewis approach to bonding

A

pairs of e- are localised in bonds or as non-bonding “lone pairs” on atoms

each bond is formed by a pair of e- shared by 2 atoms

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4
Q

octet rule

A

most main group atoms will tend to end up with ns2 np6 electronic configuration

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5
Q

importance of Lewis model

A

1st structure to describe bonding in structures of diatomics/polyatomics

introduced ideal of e- being shared

introduced bond order

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6
Q

problems with Lewis model

A

can’t easily explain radicals

difficult to explain expansion of octet

can’t explain why O2 = paramagnetic

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7
Q

paramagnetism

A

contains unpaired e-

weakly attracted by externally applied magnetic field

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8
Q

valence bond theory (VBT)

A

localised quantum mechanical approach to describe bonding in molecules

provides mathematical justification for Lewis model of e- pairs making bonds between atoms

thinks of bonds formed between 2 e- waves (bond will form if there is sufficient overlap of appropriate orbitals on 2 atoms + these orbitals are populated by a max. of 2 e-)

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9
Q

sigma bond

A

symmetrical about internuclear axis

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10
Q

pi bond

A

have node on internuclear axis + sign changes across axis

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11
Q

double bond

A

σ bond and π bond

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12
Q

triple bond

A

σ bond and 2 π bonds

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13
Q

diamagnetism

A

all e- paired

repelled by magnetic field

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14
Q

formation of π bond

A

2px orbitals interact to form a π-bonding orbital (lies above and below plane)

2py orbitals interact in a similar way, giving a π-bonding orbital that lies at 90° to px

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15
Q

bond length and bond order

A

higher bond order = shorter bond length = higher bond energy/stronger bond

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16
Q

sp3 hybridisation

A

occurs in molecules with tetrahedral shape

exhibits sp3 hybridisation on central atom

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17
Q

sp2 hybridisation

A

occurs in molecules with trigonal planar configuration

hybridised orbitals = σ bonds

unhybridised orbitals = π bonds

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18
Q

sp hybridisation

A

linear geometry

[examples]
BeH2
CO2 - sp on C and sp2 on O

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19
Q

PF5 hybridisation

A

exceeds octet

nS and nP only provide 4 - need nd orbitals (ndz2) = 5 sp3dz2

2 sets - x2 axial (180 degrees from each other)
+ x3 equatorial (120 from each other and 90 from axial bonds)

*not equivalent

20
Q

octahedral geometry

A

d2sp3 hybridisation - 6 equivalent bonds 90 degrees from each other

21
Q

MOs

A

orbitals that encompass entire molecule (constructed from AOs)

22
Q

basic rules for molecular orbital theory

A

nAOs = nMOs

to combine, AOs must have appropriate symmetry and energy

each MO must be normal and orthogonal (90) to every other MO

MO formed from weighted sums of AOs

in-phase = adding atoms
out-of-phase = subtracting atoms

23
Q

energy of bonding and anti-bonding MO

A

bonding MO = lower in energy = more stable (increased e- density between overlap due to overlap reduces repulsion)

anti-bonding MO = higher in energy = less stable (decreased e- density between overlap due to overlap increases repulsion)

24
Q

forces in anti-bonding orbital

A

attractive forces outweighed by repulsive forces

net of raising energy (reduced stability) relative to 2 atoms being apart

25
forces in chemical bond
e- held between 2 nuclei attractive forces outweigh repulsive forces net reduced energy (higher stability) relative to 2 atoms being apart
26
bond dissociation energy
net of lowering energy
27
bond order = 0
molecule doesn't exist
28
which p orbitals can't interact?
px and py/pz - don't have correct symmetry
29
which has less overlap - π or σ bonds?
π (as they're 90 from each other)
30
relative energies for bonding
σ2p < π2p < π*2p < σ*2p π = more sensitive to distance
31
isoelectronic
species with same no. of e- and same bond order
32
what happens to bond order as bond length increases?
decreases
33
sp mixing
energy difference between s and p = so small that they are close enough in energy to mix occurs in boron, carbon and nitrogen
34
evidence for sp mixing
normal MO diagram for Boron would indicate that it's diamagnetic = wrong
35
sp mixing - degree of interaction
similar energy = large interaction large energy difference = small interaction separation between nS and nP orbitals increases as atomic number increases
36
reason for stabilisation of s and p orbitals across period (left to right)
1 element to right = addition of proton in nucleus and additional e- increases nuclear charge pulls orbitals closer to nucleus but increased electron repulsion pushes them further out increased attraction = more important since e- with same quantum number don't shield each other from nuclear charge s orbital = stabilised to greater extent than p with increased z as it's more penetrating Zeff increases across row
37
penetration
2p e- = more effectively shielded than 2s 2s has radial node (2p doesn't) - presence of node means there's an area of e- density relatively close to nucleus for 2s e- 2s penetrates 1s e- ∴ 2s has higher Zeff than 2p 2s = more stabilised *e- which experience greater penetration experience less shielding ∴ increased Zeff but shield other e- more effecitvely
38
consequences of sp mixing
π = same σ = moves higher in energy (less bonding)
39
why 2s orbitals is at lower energy than 2p for lithium whereas 2s + 2p orbitals are at the same energy for Li2+ ?
[Li] 1s orbital = occupied (shields 2s/2s from Z) because of radial node, 2s penetrates 1s e- density ∴ lower in energy [Li2+] only contains a single e- no shielding
40
evidence for MO diagrams
[photonelectron spectroscopy] to construct accurate MO diagrams, we need to determine energies for individual MOs IE measured using Photoelectron Spectroscopy (PSE) sample is bombarded with photons and Eke of expelled e- is measured
41
MO diagram for heteronuclear diatomic molecules (XY)
bonding MO - more electronegative atom contributes more to bonding orbital antibonding MO - more electropositive atom contributes more to anti-bonding orbital *more electronegative atom sits lower in energy
42
when do you get sp mixing for hetereonuclear diatomic molecules?
2s AOs will only interact if atoms are next to each other in the periodic table
43
which is more stable - cyclic or linear arrangement of H3+?
although cyclic has a b.o. of 1/3 and linear = 1/2, cyclic is more stable due to lower energy in bonding MO
44
what are Walsh diagrams used for?
to estimate electronic configuration that will provide lowest overall energy + be more stable by changing shape
45
why are d orbitals not always used in hybridisation
v. high in energy compared in relation to s/p orbitals