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
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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
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Radical distribution function
How the probability of finding the electron varies with distance from the nucleus
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Node
The position where the function equals 0 (0 probability of electron being found there)
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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.
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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
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Spin correlation
Parallel spins will stay further away from each other so repulsion is reduced
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The Lewis Model
Covalent bonding occurs when valence electrons are shared between 2 atoms
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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
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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
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VSEPR theory
A multiple bond is treated as though it was a single electron pair => no resonance structure,
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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
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Stochiometry
Measures amount of substanced consumed or produced in a reaction
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Valence bond theory
Half-filled atomic orbitals on 2 atoms overlap to create a bond containing paired electrons
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How do electrons from s and p orbitals form equivalent covalent bonds?
Different atomic orbitals combine to form hybrid orbitals
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Waves can combine _______ and ________
Constructively and destructively
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Sigma bonds
Lie along the line between the atoms. They're symmetric to rotation about the bond.
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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.
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Strengths of valence bond theory
Localised models are useful in discussions of bond length, strength, bond force constant and some aspects of reactivity.
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Limitations of valence bond theory
Doesn't consider the electron distribution across the whole molecule, only considers single bonds
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Bonding molecular orbitals
When wave functions combine constructively and electron density between atoms is increased
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Anti-bonding molecular orbitals
When wave functions combine destructively and electron density between atoms is decreased
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In an anti-bonding MO...
Zero probability of electrons being found between atoms
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