Section A: Covalent Interactions Flashcards Preview

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Flashcards in Section A: Covalent Interactions Deck (40)
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1

Orbital

Region of space where electron is most probably located

2

Quantum numbers

Labels that describe shape and size of orbitals

3

Principal quantum number shows

The distance from the nucleus

4

The angular momentum quantum number shows

The shape of orbital

5

The magnetic quantum number shows

The orientation of the orbital

6

Wavefunctions

Calculate the shape of orbitals

7

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

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

8

Radical distribution function

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

9

Node

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

10

Pauli Exclusion Principle

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.

11

Hurd's rule of maximum multiplicity

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

12

Spin correlation

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

13

The Lewis Model

Covalent bonding occurs when valence electrons are shared between 2 atoms

14

Resonance hybrids

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

15

Hypervalent compounds

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

16

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

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

17

VSEPR theory

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

18

Limitations of Lewis model of bonding

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

Stochiometry

Measures amount of substanced consumed or produced in a reaction

20

Valence bond theory

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

21

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

Different atomic orbitals combine to form hybrid orbitals

22

Waves can combine _______ and ________

Constructively and destructively

23

Sigma bonds

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

24

Pi bonds

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