Orbitals and the periodic table Flashcards
1
Q
Give the n, l and ml numbers for the following atomic orbitals.
A
Next card.
2
Q
1s
A
n = 1
l = 0
ml = 0
3
Q
2s
A
n = 2
l = 0
ml = 0
4
Q
3s
A
n = 3
l = 0
ml = 0
5
Q
2pz
A
n = 2
l = 1
ml = 0
6
Q
2py
A
n = 2
l = 1
ml = -1
7
Q
2px
A
n = 2
l = 1
ml = +1
8
Q
3pz
A
n = 3
l = 1
ml = 0
9
Q
3px
A
n = 3
l = 1
ml = +1
10
Q
3py
A
n = 3
l = 1
ml = -1
11
Q
List the five 3d atomic orbitals and give their n, l and ml numbers on the next flashcards.
A
3dxy, 3dyz, 3dxz, 3dx^2-y^2, 3dz^2
12
Q
3dxy
A
n = 3
l = 2
ml = -2
12
Q
3dxz
A
n = 3
l = 2
ml = +1
13
Q
3dx^2-y^2
A
n = 3
l = 2
ml = +2
13
Q
3dyz
A
n = 3
l = 2
ml = -1
13
Q
3dz^2
A
n = 3
l = 2
ml = 0
13
Q
Does it matter which 2p, 3p or 3d orbital has which ml number?
A
For subshells like 2π 3π, or 3π, it doesnβt matter which orbital is assigned to which ππ
value. The assignments are interchangeable and chosen based on convention, without affecting the physical or chemical properties of the atom.
14
Q
Images of orbitals are on the summary sheet.
What do the red and blue mean?
A
Alternating wave phases.
15
Q
What does wave phase refer to?
A
Oscillating sign (+ve or -ve) of a wavefunction (πΏ).
16
Q
When πΏ>0, what is the phase?
(refer to summary sheet for a full diagram of a wave and +ve/-ve half cycles)
A
Positive.
17
Q
What is a wavefunction?
A
Describes the probability distribution of electrons around an atom.
18
Q
When πΏ<0, what is the phase?
A
Negative
19
Q
n, l, ml and ms are the allowed eigenvalues for atomic orbital wavefunction.
What is n?
Give the range of numbers it can be.
A
Principle quantum number.
Can take any positive integer.
20
Q
What is l?
Give the range of numbers it can be.
A
Secondary/angular/azimuthal/orbital quantum number.
0 to (n-1).
20
What is ms?
Give the range of numbers it can be.
Spin quantum number.
-1/2 or +1/2
(for spin down or spin up)
21
What is the Aufbau principle?
Electrons fill subshells of the lowest available energy in the ground state of an atom/ion.
21
What is ml?
Give the range of numbers it can be.
Magnetic quantum number.
-l to +l
22
What is Hund's rule of maximum multiplicity?
In the lowest energy electron configuration, electrons fill orbitals singly before pairing (they have parallel spins).
23
What is the Pauli exclusion principle?
No two electrons can have exactly the same set of four quantum numbers: n, l, ml, ms.
24
Give some examples of hydrogen (one electron) bodies.
H, He+, Li2+
25
Give 3D Cartesian co-ordinates.
x, y, z
26
Give 3D polar co-ordinates.
ΞΈ, Ο, r
(angle, elevation, distance)
27
Write 3 sets of co-ordinates to illustrate how electrons are represented as a wave function.
πΏ(x, y, z)
πΏ(ΞΈ, Ο, r)
πΏ(n, l, ml)
28
Write the wavefunction broken into two smaller functions.
Explain what they are.
R(n) = radial wave function
- describes how the probability density of
an electron varies with distance from the
nucleus in an atom.
Y(ml, l) = angular wave function
- describes the orientation and shape of the
electron's probability distribution
29
What is effective nuclear charge?
The effective nuclear charge (πeff) is the net positive charge experienced by an electron in a multi-electron atom. It represents the actual attractive force that a specific electron "feels" from the nucleus, after accounting for the shielding effects of other electrons.
The effective nuclear charge experienced by an electron is typically calculated as:
πeff=πβπ
where π is the atomic number (total positive charge from the protons), and π is the shielding constant, which represents the average repulsion effect from other electrons.
Slater's rules = method to estimate shielding constant.
30
When the radial distribution function (4Οr^2R(r)^2) is plotted against R(r)^2, what does this show?
Why are there no negative values?
(refer to small diagram in summary sheet and full diagram in course manual)
The probability of finding an electron at distance r from the nucleus.
R(r)^2
31
There is a peak in the 1s curve. What does this tell us?
The most probable distance from the nucleus for 1s electrons.
This is equivalent to the Bohr radius in the Bohr model.
32
What is a radial node?
The probability of finding an electron at a particular distance from the nucleus is 0 because the electron density drops to 0.
33
Why is the 1s peak shifted closer to the y axis than the 2s peak?
The majority of electron density of 1s is closer to the nucleus than 2s (there is an overlap).
34
How do you work out how many radial nodes an s orbital peak has?
n - 1
35
How do you work out how many radial nodes a p orbital has for r>0?
n - 2
36
How do you work out how many radial nodes a d orbital has for r>0?
n - 3
37
How do you work out how many radial nodes an f orbital have for r>0
n - 4
38
Generally, for r>0, the first time an orbital occurs it has no...
radial node (1s, 2p, 3d, 4f orbitals).
39
As n increases, the peaks of a radial wave functions become...
broad and shallow.
40
Why is this?
Due to the increase in the size and spread of orbitals, formation of additional radial nodes and decreased electron density near the nucleus.
41
Describe and explain the periodic trend in atomic radii (pm).
(see the course manual for all these periodic trend graphs)
Graph = atomic radii as a function of atomic number.
Generally atomic radii decrease across a period due to increasing nuclear charge so electrons are pulled closer to the nucleus.
Generally atomic radii increase down a group due to more electron-electron repulsion between valence and core electrons and increasing quantum number (so constant effective nuclear charge).
42
What are the exceptions to atomic radii trends?
Exceptions in K-Br and Rb-I on graphs are due to electrons moving into orbitals for electron-electron pairing.
Exceptions from Cs-P due to d-orbitals experiencing lots of penetration. D
43
Define metallic/covalent radius.
1/2 the distance between nuclei of adjacent atoms.
44
What kind of atomic radius do group 18 elements have? Why?
Van der Waals - they are monatomic.
45
What is a Van der Waals radius?
1/2 distance between nuclei of two identical non-bonded atoms.
46
Give the periodic trends in ionisation energy (kJ/mol).
Graph = first ionisation energy as a function of atomic number.
Across period first ionisation energy increases due to increasing nuclear charge, similar shielding and decreasing atomic radius.
Down group first ionisation energy decreases due to increasing atomic radius and shielding which decreases effective nuclear charge.
47
Give exceptions to the periodic trends in ionisation energy.
Very high first ionisation energies for group 18 due to stable electron configuration (full valence shells).
Exceptions due to outermost electron being in higher energy orbitals so easier to remove and electrons being paired so easier to remove because of repulsion.
48
Give the periodic trends of first electron affinity (kJ/mol).
Graph = electron affinity as a function of atomic number.
Across period first electron affinity becomes more negative due to increasing nuclear charge (more favourable to accept electrons), decreasing atomic radius (so greater effective nuclear charge) and elements closer to filling p orbitals want to gain electrons.
Down group, becomes less negative due to increasing atomic radius and more shielding.
49
What are the exceptions to the first electron affinity trends?
Group 2 - less negative than expected due to full s subshell.
Group 15 - less negative than expected due to electron-electron repulsion caused by adding electrons (e.g. N = 2p3).
Noble gases - positive values as they don't readily gain electrons.
F slightly less negative than Cl due to small atom causing electron-electron repulsion.
