Atomic Orbitals And Electronic Configuration Flashcards
What’s the definition of orbital
An orbital tells us the probability of finding an electron in a certain region of space
How many electrons can each orbital hold?
A maximum of two
How many types do orbitals are there? How do they differ?
Four - s, p, d, and f-orbitals
They differ in their shape and energy
S orbitals:
Shape?
Total number of electrons?
Spherical shape
2 electrons in total
P orbitals:
Number of orbitals?
Shape?
Number of total electrons?
X3 orbitals
Dumbbell shape
6 electrons in total
D orbitals:
Number of orbitals?
Number of total electrons?
X5 orbitals
10 electrons in total
F orbitals:
Number of orbitals?
Number of total electrons
X7 orbitals
14 electrons in total
Shell number: 1
No. of orbitals in the shell?
Names of orbitals in the shell?
Max No. of electrons?
1, S, 2
Shell number: 2
No. of orbitals in the shell?
Names of orbitals in the shell?
Max No. of electrons?
2
S, P
8
Shell number: 3
No. of orbitals in the shell?
Names of orbitals in the shell?
Max No. of electrons?
3
S, P, D
18
Shell number: 4
No. of orbitals in the shell?
Names of orbitals in the shell?
Max No. of electrons?
4
S, P, D, F
32
Rule 1 of orbitals:
The aufbau principle
The way we fill up orbitals
Electrons in an atom occupy the lowest possible energy levels or orbitals first
What is the ‘boxes in a line’ method?
The electrons are represented by arrows which occupy the boxes. One arrow points upwards the other down.
Why do the arrows face opposite ways in the box method?
This is known as spin pairing and occurs to minimise the
repulsions between electrons.
Why is shell 1 always filled first?
They are lowest in energy
Rule 2 of orbitals:
Pauli Exclusion Principle
No orbitals can contain more than two electrons and when two electrons occupy the same orbital they do so spinning in the opposite direction
Orbital rule 3:
Hund’s rule
Every orbital within a subshell must fill up singly before pairing the electrons
What do you do after the orbitals up to 3p have been filled? Why?
You fill 4s before 3d, then go back before moving on. This is because the 4s orbital is lower in energy than the 3d
What’s a quicker way of writing electronic configurations?
Short hand notation
Leave out the boxes and abbreviate
E.g.
Lithium = 1s^2 2s^1
Hydrogen = 1s^1
1 = energy level (principle quantum number)
S = number of electrons in orbital
^1 = type of orbital
Exceptions to the aufbau principle? Why?
Cu and Cr are anomalies because they only have 1 electron in the 4s
Because the 4s and 3d orbitals are close in energy it is not that difficult for one of the 4s electrons to “hop” from the 4s to the 3d (it also has to change its spin).
There is a measure of stability associated with full and half-full sets of orbitals.
This is much more favourable situation since now both the 4s and the 3d orbital sets are both half full. Admittedly the 4s is no longer completely full but this is compensated for by the gained stability of the now half-full 3d set. A similar reasoning can be applied to explain the configuration of copper.
What do you do when writing the electronic configuration of ions?
Ions are atoms which have either lost or gained electrons. Therefore in writing the electronic configurations of ions you have to remember to add on the electrons gained, or take away the electrons lost from the outer orbital first.
When writing the configuration for transition metal ions, do you remove the electrons from 4s orbitals before 3d or not?
You do
This is because the 4s orbital is furthest from the nucleus (and screened by inner 3d orbital’s), so electrons are lost from here before the 3d orbital’s.
(The actual explanation as to why this is, is actually quite complex. If you wish to read more on this, see AS chemistry moodle)
