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Chemistry IA > Week 3 Module 1 > Flashcards

Week 3 Module 1 Flashcards

1
Q

What is the lowest allowed energy of the Hydrogen Atom?

A

E1=-2.18x10^-18J (for n=1), then E2=-5.45x10^-19K (for n=2)

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2
Q

What are molecular orbitals composed of?

A

Atomic orbitals. The orbitals can overlap and be different depending on the phase. Relates to chemical bonding.

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3
Q

What happens to energy as n increases?

A

En approaches the energy of an unbound electron - 0.

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4
Q

How many types of orbitals are there with n=2?

A

Two types: 2s and 2p.
4 orbitals share the n=3 energy - (there is 3 types of 2p)

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5
Q

What are two orbitals that share the same energy called?

A

Degenerate orbitals.
eg. 2s and 2p.

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6
Q

What does energy depend on?

A

It only depends on n as E(R) is a constant. Energy does not care about l or m, so the difference doesnt matter. Same for any atoms with 1 electron.

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7
Q

What happens if you remove 1 electron from helium?

A

It becomes helium+, a hydrogen-like atom.
Seen with Lithium2+ etc.

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8
Q

What is the equation for the allowed energies of Hydrogen-Like atoms?

A

En=-E(R) x (z^2/n^2)

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9
Q

What does z represent in En=-E(R) x (z^2/n^2)?

A

The atomic number of the atom being handled. Eg. for helium it would result in 4 and lithium would be 9.

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10
Q

What would happen if I gave my n=1 energy electron more energy and it quantum transitioned into n=2?

A

It would not be happy as it now has (pi?) energy it wants to give off to emit light. As a result, it gives off energy and goes through quantum transition back.

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11
Q

What is the Hydrogen spectrum?

A

What causes specific wavelength colours depending on n value. Specific lines correspond to hydrogen - can be seen within the Sun.

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12
Q

Where are Quantum transitions seen in evergy day life?

A

Neon light in a neon lamp given energy.

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13
Q

What can we do by looking at the intensity of quantum transitions in a star?

A

We can tell the temperature.

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14
Q

When does a energy transition happen?

A

When energy is released or absorbed from a particular energy level.

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15
Q

How is ΔE (difference between energy levels) calculated?

A

E(n2) - E(n1) = -Z^2 x E(R) x ((1/n2^2) - (1/n1^2)
It works for H or H-like atoms.

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16
Q

How can energy be provided to cause a quantum transition?

A

An electrical current, heat, etc.

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17
Q

What is the equation for E(n)?

A

-(2.18 x 10^-18J)/n^2 (h?)

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18
Q

What occurs if you provide even more energy to an atom?

A

There is an even greater number of possible transitions.

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19
Q

What region do transitions that involve really high energy belong in?

A

Ultraviolet region as our eyes cannot see them. That is why Hydrogen isnt use in a lamp.
6 -> 2 is violet, 5 -> 2 purple, 4 -> 2 cyan, 3 ->2 red, etc.

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20
Q

Worked example of Hydrogen that we want to use a photon to allow n=1 to n=3. We want to calculate the wavelength of the light in 3 sig figs.

A

ΔE = E(3) - E(1)
= -Z^2 x E(R) ((1/n(2)^2) - (1/n(1)^2))
= - 1^2 x (2.18 x 10^-18) x (1/3^3 - 1/1^2)
= 1.94 x 10 ^ -18 J
As we are looking for wavelength: ΔE = hc / λ
Therefore, λ = hc / ΔE
= 1.02 x 10^-7m

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21
Q

How can atomic orbitals vary?

A

They come in many shapes and sizes and different energy (which depends on n)

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22
Q

What must we think of when we consider molecules?

A

As they are made of atoms we must think of the atoms first.

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23
Q

How are atomic orbitals connected to molecular orbitals?

A

They are just 3D waves which can be used to make molecular orbitals.
Molecular orbitals also have shapes, lobes, and nodes,as well as electron density and energy.

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24
Q

How can we think of a H2 molecule?

A

We can think of a system where we have 2 protons and 2 electrons - how far are these atoms seperated? (the length of the chemical bond)
More complicated as there are several atoms involved.

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25
What can calculations about a H2 molecule answer?
- Whether a bond forms - Equilibrium bond length - Structure of the bond - Electronic properties of the molecule (eg. transitions)
26
What does Molecular Orbital Theory (MOT) solve?
The problem of multiple electrons.
27
What will happen if there is 2 1s orbitals in a molecule (eg. H2)?
They will interfere in one of two ways. Each molecular orbital contains a max of two (spin paired) electrons.
28
Why can bonds form between two atoms in a molecule?
Because it can cause the electrons to overall have a lower energy than when the electrons are seperated.
29
In what situation would a molecule be unstable?
If the energy is hgher or the same when 2 atoms are put together compared to when they are apart, making there be no purpose in the molecule.
30
What happens to electrons in a molecule? (maybe in all molecules, im not sure)
The electrons are delocalised, spreading all over the molecule. Electrons now are attracted more strongly as there is two nuclei.
31
What forms a sigma 1s orbital?
Formed by 2 1s orbitals in the same phase giving constructive interference/overlap (opposed to destructive interference). No longer spherical but now oblong. It is why wehave the H2 molecule.
32
What happens when 2 s orbitals overlap?
They form a σ-orbital. H2 molcule description is σ1s^2 (2 x exponent - 2 electrons)
33
What orbital is formed from destructive overlap?
The anti-bonding orbital σ1s^* Higher energy compared to constructive. No amplitude between, therefore there is a node.
34
Compare σ1s and σ*1s to 1s orbital in energy.
σ1s has lower energy than the 1s orbital while σ*1s is higher. The differencer btween σ1s and 1s is the same difference as 1s and σ*1s.
35
Is energy gained forming molecular orbitals?
No (conservation of energy laws), but it depends on the electrons. H2 has 2 electrons that prefer to pair up on the lowest energy orbital. Both electrons go into σ1s instead of σ*1s and there is a driving force to become a molecule - its stable.
36
What is a bond orderl?
The quantity we can use to see if a molecule is stable. The bond order (single, double, triple, etc) of a covalent bond can be found by molecule orbital diagram: Bond Order = (1/2) x [(# of bonding electrons) - (# of anti-bonding electrons)]
37
What is the bond order of H2 and what does it mean?
H2 is (1/2) x (2 - 0) = 1. B.O. of 1 means it is stable. It can go up to 3. Unstable is 0.
38
What does a higher bond order mean.
The higher the bond order, the higher the bond strength/energy.
39
What is the relationship between bond order and bond distance?
Inversely proportional.
40
Explain He2
Noble gas, unreactive. He has 2 electrons, so He2 has 4. The first two electrons go into σ1s, the second two go into σ*1s. As a result, B.O. = 0. Same number in bonding and anti-bonding, therefore it is unstable and there is no bond. He is very stable by itself.
41
What are the types of diatomic molecules in CHEM1A?
1. Homonuclear Diatomics: Formed from 2 identical atoms. eg. F2 2. Heteronuclear Diatomics: Formed from 2 different atoms. Eg. ClBr. - we focus on homonuclear.
42
What is the bond length in homonuclear diatomic molecules used to define?
The "Covalent radius" of the atom.
43
What happens with energy of atomic orbitals when >1 electron in atom.
Energy levels with same n but different l (s,p,d) are no longer degenerate.
44
Compare the energy of 2s orbitals to 2p in a multi-electron atom.
2s orbitals have slightly less energy than 2p due to it being after 1s and 2s are already filled.
45
Compare energy of 3s, 3p, and 3d orbitals.
In H atom, they would have the same energy. When there are multiple atoms, the energy increases. (similar reasoning to 2s and 2p).
46
Which has higher energy in a multi electron atom: 4s or 3d?
3d has higher energy than 4s.
47
Explain Li2.
Li as 3 electrons. therefore Li2 has 6 electrons.We fill them into the orbitals, lowering of energy happens with σ2s orbital and thus formation of the σ*2s. σ1s anf σ*1s are still there, and then2 electrons are added into the σ2s orbital. Therefore, B.O. = (1/2) x (4-2) = 1 Li2 is stable.
48
Explain Be2.
Be has 4 electrons, Be2 has 8. 2 in σ1s, σ*1s, σ2s, σ*2s. B.O. = (4-4) = 0 Be2 is unstable.
49
Compare the energy of σ1s and σ*1s to sigma σ2s and σ*2s.
σ1s and σ*1s are much lower.
50
What is the shape of 2p orbitals?
The shape of an 8. There are two phases. When there is overlap, the phases are important, forming a σ-type molecular orbital. (refer to page 19)
51
What happens with 2p orbitals that have the same phase in the middle?
There is constructive overlap and they bond in the middle. The two side lobes stay the same. Note: ψ has max value between the nuclei and 2 planar nodes
52
What happens with 2p orbitals that don't have the same phase in the middle?
There is destructive overlap and σ* anti orbital is produced. ψ has 0 value between nuclei, and 3 planar nodes. - Higher in energy than σ2p
53
Explain π orbitals from the interaction of p atomic orbitals.
2p atomic orbitals that have side by side overlap form a π-orbital. ψ is dominant above and below the bond but between the nuclei - it is a type of covalent bond. 2 nodes merger - only 1 nodal plane.
54
Explain π* orbitals from the interaction of p atomic orbitals.
ψ has 0 value between the nuclei. Anti-bonding. 2 orthogonal nodal planes. Destructive overlap. Y-axis is very similar to the x-axis.
55
What does overlap of p(y) orbitals produce?
Identical π and π* molecular orbitals that point out of the plane. πp(x) and πp(y) are degenerate. π*p(x) and π*p(y) are also degenerate.
56
Compare πp(x) to πp(y).
Identical but different orientation (refer to page 20)
57
Relative energies of molecular orbitals?
Basically ranked by nodes. (try find reference to any week3 lecture 2)
58
Explain B2.
B has 5 electrons, B2 has 10. 9 and 10 are unpaired e-s in π(2px) and π(2py) in the same direction (same spin). Same spin is lower energy (favourable). B.O. = (1/2) x (6-4) = 1. When you take B2 and bring it to a magnet, it would interact due to the unpaired electrons.
59
Explain C2.
Present in space - front part of a comet is green cause of C2. C has 6 electrons. C2:12. 11 and 12 are added to π(2px) and π(2py) with the unpaired electrons of the same spin. It is not paramagnetic - it is diamagnetic (would not interact with magnets). B.O. = (1/2) x (8-4) = 2
60
Explain N2.
Like C2 but with 2 extra electrons that fill σ(2pz) B.O. = (1/2) x (10-4) = 3. Very very stable, highest order we'll see in this discussion.
61
Explain isomers in relation to bonds.
One bond points up, other points down. In some cases, a particular isomer is preferred. Comes out of orbital overlap. If orbital overlap results in a bond that points up, that will be majority of isomer. Nature works taht way with orbital overlap influencing synthesis.
62
What makes a molecule paramagnetic?
Unpaired electrons. This is why although O2 appears diamagnetic, it is paramagnetic. (memorise order of orbitals sigma to pi etc)
63
What would happen if you removed an electron from N2?
It would become N2^+ and become paramagnetic.
64
What makes O2 and F2 different?
σ2p and π2p are flipped with σ2p being after σ2p*. (do not need to mnow why).
65
Explain O2.
B.O. = (1/2) x (10 - 6) = 2. Due to lone electrons in π2px* and π2py*, it is paramagnetic. Lower than N2 bond order. π being lower - increase in antibonding electrons.
66
What is the secret evil 4th quantum number?
M(s) <- the spin quantum number.
67
What do all 4 quantum numbers allow us to identify?
Any electrons within any atomic orbital.
68
How does spin work with electrons?
Electrons can have a spin that is spin up or spin down. If there are two, there must be 1 spin up and 1 spin down. Spin up spin up would take an infinite amount of energy to occur.
69
How is spin up represented by M(s)
M(s) = +(1/2) Spin down is -(1/2) Paired electrons give a total spin of 0, and exibit no magnetism - are diamagnetic.
70
Are most compounds diamagnetic or paramagnetic?
Most are diamagnetic.
71
What happens when unpaired electrons point in the same direction?
It givesthe molecule a magnetic property - they become paramagnetic. Tiny magnetic fields are created by the electrons.
72
What do you need to remember when filling out electron shell chart?
To prioritise unpaired electrons first.

Chemistry IA (6 decks)

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