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2X Chemistry > Quantum Mechanics > Flashcards

Quantum Mechanics Flashcards

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1
Q

what does the schrodinger equation explain

A

how particles (electrons) behave

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2
Q

what is the quantum object characterised by

A

a wave function

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3
Q

whats the probability that a particle will hope to a different place given by

A

a quantity called the action

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4
Q

explain wave function duality

A

-objects bigger than a molecule have a negligible wavelength
-an electron is tiny so its wavelength is around the size of an atom
-electrons are viewed as particles+waves

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5
Q

how would you describe electrons in terms of waves in an atom

A

they’re circular standing waves surrounding the nucleus

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6
Q

what is an objects wavelength inversely proportional to

A

its mass

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7
Q

how do you calculate the energy of a photon

A

=planks constant x frequency
=h x f

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8
Q

explain what the wave function is

A

-a probability amplitude
-the square of the magnitude of the wave function describes the probability of an electron existing in an exact location

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9
Q

explain an electron is terms of quantum mechanics

A

a cloud of probability density

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10
Q

summarise the main info about schrodinger equation

A

-applies the quantum systems
-describes the systems 3-dimensional wave function
-calculates the wave function of a system
-the info it contains is probabilistic (Heisenberg uncertainty principle)

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11
Q

what does the Schrodinger equation tell us

A

where/how the particle is inside the walls

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12
Q

what is assumed for a particle in a 1D box

A

a potential V(x)=infinity, everywhere except in between x=0 and x=a, where V(x)=0

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13
Q

what does it mean when V(x)=infinity

A

a mathematical way of saying the particle is definitely not there

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14
Q

what does P(x) when V(x)=infintity

A

P(x)=0
the probability density function equals zero

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15
Q

what 3 quantum numbers are acceptable wave functions numbered by

A

n=principle quantum number
l=orbital angular momentum quantum number
MI=magnetic orbital quantum numbers

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16
Q

what does n describe

A

size +energy of that orbital

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17
Q

what does l describe

A

shape of the orbital

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18
Q

what does MI describe

A

orientation of the orbital

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19
Q

list the relationship between between these 3 quantum numbers

A

N=1,2,3
L=0…n-1 e.g. when n=3, l=0,1,2
MI= -l…l e.g. when l=1, MI=-1,0,1

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20
Q

what is the wave function for an electron called

A

an orbital

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21
Q

describe the quantum numbers of a degenerate orbital

A

-the same energy, n
-different l and MI

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22
Q

for the h atom describe the orbitals

A

all n^2 orbitals for a given value of n are degenerate

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23
Q

what does the radial part (r)determine

A

the spatial extent of the wavefunction

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24
Q

how do you calculate how many nodes an atomic orbital has

A

=n-1

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25
Q

what does the radial distribution function give

A

the probability of finding the electron in spherical shell of thickness at distance r

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26
Q

what does the angular part of an orbital determine

A

the shape of the orbital

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26
Q

what does ref equal at the nucleus

A

0

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27
Q

what is an orbital a product of

A

=radial part x angular part

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28
Q

what the angular part of an s orbital

A

-they have no angular dependance
-they’re spherically symmetric

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29
Q

compare the radial part of 2s in comparison to 2p

A

radial part of 2s is more highly curved(1 radial node)than radial part of 2p(no radial nodes) so 2s has more radial kinetic energy so the e- has more energy moving in and out of the nucleus

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30
Q

compare the angular part of 2s in comparison to 2p

A

angular part of 2s(no angular node) is less curved than angular part of 2p(1 angular node), 2p has more angular kinetic energy so the energy is used moving around the nucleus

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31
Q

what is an atomic orbitals nodes made up of

A

radial +angular nodes

32
Q

list how many +what type of nodes a 1s,2s,2p,3s,3p atomic orbitals have

A

1s=0 nodes
2s=1 node(radial)
2p=1 node(angular)
3s=2 nodes(both radial)
3p=2 nodes(1 radial+1 angular)

33
Q

what is angular momentum quantised by

A

magnitude+length of vector

34
Q

why is the wavefunction higher closer to the nucleus

A

-as the electron is attracted to a more positive nucleus
-so there’s a contraction of orbitals
-so the electrons now have lower energy so it’s a more stable system

35
Q

why does the SE only give approximate solutions for atoms with more than 1 molecule

A

as the electrons interact with each other making it too complicated

36
Q

whats the relationship between angular momentum,l and magnetic dipole moment, u

A

they’re always anti-parallel to each other

37
Q

what type of spin do particles have

A

-integer spin
-half-integer spin

38
Q

what’s an electrons spin quantum number

A

s=1/2

39
Q

what’s the magnetic spin quantum number(ms) for an electron

A

-s…+s

40
Q

what values are magnetic spin quantum numbers restricted to

A

2s+1

41
Q

what’s another word for spin-up

A

alpha

42
Q

what’s another word for spin-down

A

beta

43
Q

define Zeff

A

the effective nuclear charge, the net positive charge pulling these electrons towards the nucleus

44
Q

why does the 2s orbital penetrate the 1s2 orbital

A

because the 2s orbital with its inner radial node has a inner peak of electron density that partially penetrates the shielding of the 1s2 electrons

45
Q

why is Zeff lower for 2p than the 2s orbital

A

as the 2p orbital doesn’t have an inner radial node so doesn’t have an inner peak of electron density+so has greater shielding from 1s2 electrons

46
Q

why is the 2s orbital lower in energy than the 2p orbital

A

as it has a higher Zeff so a stronger attraction

47
Q

what is shielding

A

when inner electrons partially shield outer electrons from full Z

48
Q

what’s penetration

A

when an e- in an outer orbital has a significant probability of being found inside an inner orbital

49
Q

what type of orbitals are 2s+2p for any 1 electron ion

A

degenerate

50
Q

what happens to atomic radius and down a group and across a group

A

-increases down a group
-decreases across a period

51
Q

why does atomic radius increase down a group

A

-n increases by 1
-so the number of completely filled shells increases
-these shield the nuclear charge

52
Q

why does atomic radius decrease across a period

A

-increasing nuclear charge
-so electrons are attracted more strongly
-atom is more compact

53
Q

define first ionisation energy

A

energy required to remove 1 electron from the neutral atom in the gas phase

54
Q

what happens to ionisation energy down a group and across a period

A

-increases down a group
-decreases across a period

55
Q

why does ionisation energy decrease down a group

A

-completed inner shells of electrons shield
-valence electron is attracted less to the nucleus

56
Q

why does ionisation energy increase across a period

A

-increasing nuclear charge
-electrons are bound more strongly

57
Q

which elements don’t follow the average trend for ionisation energies

A

Li
B
O

58
Q

why odes Li not follow the ionisation energy trend

A

as there’s an e- in the 2s orbital, which is well shielded by the completed (1s)^2 shell

59
Q

why does B not follow the average ionisation trend

A

it’s the first time the 2p is occupied, p electrons are less strongly bound than s electrons

60
Q

why does O not follow the average ionisation trend

A

as its the first time any of the 2p orbitals are doubly occupied, e-e repulsion is higher for an electron in a doubly occupied orbital, lowering the binding energy

61
Q

what does HOMO stand for

A

highest occupied molecular orbital

62
Q

what does LOMO stand for

A

lowest unoccupied molecular orbital

63
Q

what does UV absorption occur from

A

excitation of an electron from filled to empty orbital

64
Q

what’s a destructive wave

A

when a positive+ negative wave cancel each other out so the resultant becomes zero

65
Q

what happens when you combine separate atoms in a molecule

A

constructive interference

66
Q

what happens when constructive interference occurs

A

-overlap of wave functions
-results in increase in value of wavefunction
-increase in electron density

67
Q

describe an antibonding combination

A

-positive+negative wave functions added together results in close to zero wave function
-destructive interference results in low probability of finding the electron
-no electron density in the region between the nucleus

68
Q

how are molecular orbitals built up

A

by adding or subtracting atomic orbitals

69
Q

describe the electrons in bonding MO

A

-high probability of being close to both nuclei, build-up of density between the nuclei
-internuclear density +nuclei attract each other, stabilising
-density pulls nuclei together, bonding
-more stable in separate atoms

70
Q

describe electrons in the anti bonding MO

A

-zero probability of being close to both nuclei, depletion of density between the nuclei
-nuclei pulled apart by density outside the internuclear region
-node means that electrons have high kinetic energy which is destabilising
-less stable than in separated atoms

71
Q

what do molecular orbital diagrams

A

connection between shape, energy and phase

72
Q

how do you calculate bond order

A

(number or e- in bonding MOs- number of e- in anti bonding MOs)/2

73
Q

how are sigma-bonding MOs formed

A

from overlapping s-orbitals or from p-orbitals oriented along the bond axis

74
Q

how are pi-bonding MOs formed

A

from overlapping p-orbitals oriented perpendicular to bond axis

75
Q

why is oxygen paramagnetic

A

due to unpaired oxygen

76
Q

why is N2 diamagnetic

A

there are no unpaired electrons

77
Q

what’s the relationship between electronegativity of an atom +its orbital energies

A

the more electronegative an atom, the lower its orbital energies

78
Q
A

2X Chemistry (3 decks)

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