Section A: Covalent Interactions Flashcards

(40 cards)

1
Q

Orbital

A

Region of space where electron is most probably located

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2
Q

Quantum numbers

A

Labels that describe shape and size of orbitals

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3
Q

Principal quantum number shows

A

The distance from the nucleus

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4
Q

The angular momentum quantum number shows

A

The shape of orbital

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5
Q

The magnetic quantum number shows

A

The orientation of the orbital

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6
Q

Wavefunctions

A

Calculate the shape of orbitals

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7
Q

How do wavefunctions explain the order of energies of the orbital?

A
  1. The further the orbital from the nucleus, the higher the energy
  2. Wavefunctions can be used to calculate radical distribution function
  3. These distribution functions can be found for all the orbitals and explain relative energies of 2p and 2s orbitals
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8
Q

Radical distribution function

A

How the probability of finding the electron varies with distance from the nucleus

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9
Q

Node

A

The position where the function equals 0 (0 probability of electron being found there)

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10
Q

Pauli Exclusion Principle

A

Electrons have spin, related to angular momentum. No 2 electrons in an atom can have the same 4 quantum numbers. So only 2 electrons can occupy an orbital.

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11
Q

Hurd’s rule of maximum multiplicity

A

If 2 or more orbitals have the same energy, electrons will spread out to occupy the maximum number of these, maximizing the number of parallel spins

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12
Q

Spin correlation

A

Parallel spins will stay further away from each other so repulsion is reduced

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13
Q

The Lewis Model

A

Covalent bonding occurs when valence electrons are shared between 2 atoms

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14
Q

Resonance hybrids

A

The true structure of a molecule can be considered an average of resonance forms

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15
Q

Hypervalent compounds

A

Compounds which require more than an octet of electrons in order to draw a Lewis structure

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16
Q

Assumptions of the valence shell electron pair repulsion theory (VSEPR)

A
  1. Electrons in bonds and lone pairs form clouds of electron density that repel one another
  2. The lowest energy arrangement is when they’re farthest apart, determining equilibrium molecular shape
  3. Lone pairs repel more than bonding pairs
  4. Lone pair prefers equatorial position over an axial position -> less repulsion
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17
Q

VSEPR theory

A

A multiple bond is treated as though it was a single electron pair => no resonance structure,

18
Q

Limitations of Lewis model of bonding

A
  1. Doesn’t explain hypervalency
  2. Doesn’t explain unpaired electrons -> paramagnetic
  3. Doesn’t explain what sharing of electrons means or how it arises
19
Q

Stochiometry

A

Measures amount of substanced consumed or produced in a reaction

20
Q

Valence bond theory

A

Half-filled atomic orbitals on 2 atoms overlap to create a bond containing paired electrons

21
Q

How do electrons from s and p orbitals form equivalent covalent bonds?

A

Different atomic orbitals combine to form hybrid orbitals

22
Q

Waves can combine _______ and ________

A

Constructively and destructively

23
Q

Sigma bonds

A

Lie along the line between the atoms. They’re symmetric to rotation about the bond.

24
Q

Pi bonds

A

They form side-ways on. Two lobes, above and below the line between the atoms. They’re not symmetric to rotation about the bond.

25
Strengths of valence bond theory
Localised models are useful in discussions of bond length, strength, bond force constant and some aspects of reactivity.
26
Limitations of valence bond theory
Doesn't consider the electron distribution across the whole molecule, only considers single bonds
27
Bonding molecular orbitals
When wave functions combine constructively and electron density between atoms is increased
28
Anti-bonding molecular orbitals
When wave functions combine destructively and electron density between atoms is decreased
29
In an anti-bonding MO...
Zero probability of electrons being found between atoms
30
What is the relative energy of MO's?
Bonding MO is lower in energy than orginial orbitals and anti-bonding MO is higher -> fill lowest energy MO first
31
If electrons occupy the same MO they must have
Opposite spin
32
Number of MO =
Number of AO
33
Only orbitals of the same ____ wil combine
Symmetry
34
Bond order =
( Number of bonding electrons - number of anti bonding electrons ) / 2
35
A bond order of 0 means
There is no bond
36
The degree of overlap of AO's is reflected in ...
The separation in energy of bonding and anti-bonding MO's (larger separation = stronger bond)
37
How are MO's formed from p orbitals?
1. Pz orbitals overlap head on -> sigma MO's | 2. Py and Px orbitals overlap side on -> pi MO's
38
Limitations of MO theory
Difficult to predict order in energy of orbitals in more complex molecules
39
Why is liquid oxygen attracted to a magnet?
1 pi anti-bonding MO each contain an unpaired electron
40
Heisenberg uncertainty principle
The position and the velocity of an object cannot both be measured exactly, at the same time