Outside his buckyball home, one molecule overheard another molecule saying, "I'm positive that a free electron once stripped me of an electron after he lepton me. You gotta keep your ion them."
Bahahaha, get it?!?! We all need some cheesy chemistry jokes once in a while :-)
Moving on now...
We had a very perplexing class today. Brace yourselves, for we will be covering a difficult topic: ELECTRONIC CONFIGURATION. So pay attention carefully!
What is it?
Electronic configuration of an atom is a notation that describe the orbitals where the electrons occupy and the total number of the electrons in each orbital.
Neils Bohr, the same guy that invented the Bohr diagram, proposed that electrons only exist in specific energy states. Electrons can move from one orbital to another when they absorb or emit a specific amount of energy.
Energy levels are the amount of energy that an electron can have. The lowest energy level is in the innermost orbit that forms a shell around the nucleus. Each orbit is assigned a quantum number (n) which indicates the number of the energy level. As n increases, the energy of the electron increases. The energy difference between two particular energy levels is called the quantum of energy.
Atoms with all electrons in the lowest energy levels are in a ground state. When electrons move up in the energy levels, they become less stable and more "excited", thus are in an excited state.
As mentioned before, electronic configuration revolves around orbitals. So what are they? An orbital is the actual region of space occupied by an electron in a particular energy level. The set of all orbitals having the same n-value or energy level is the shell. A subshell is a set of orbitals of the same type.
There are many different types of orbitals and each have their own fundamental shape. They are designated by letters s, p, d,and f. Each orbit holds a maximum of 2 electrons. (Pauli Exclusion Principle)
Specific types of orbitals are possible for a given value of "n".
n=1 : s-type ONLY
n=2 : s- and p-types
n=3 : s- , p- , and d-types
n=4 : s- , p- , d- and f-types
S-type subshell
- consists of 1 s-orbital
- has a spherical shape
- max # of electrons: 2
P-type subshell
- consists of 3 p-orbitals
- shaped like a dumbbell
- max # of electrons: 6
D-type subshell
- consists of 5 orbitals
- max # of electrons: 10
F-type subshell
- consists of 7 f-orbitals
- max # of electrons: 14
____________________________________________________________________________
Writing Electronic Configurations
(for neutral atoms)
STEPS:
- Figure out how many electrons you have (if neutral atom, just look at the atomic number)
- Follow the direction of the arrows and start with the lowest energy level first (Aufbau principle)
- Starting from 1s, keep adding the notations until you get to the designated number (of electrons). Note that each orbital has a diff maximum of electrons! (eg. s-types can only go up to 2)
Each electron has an opposite spin designated by upward and downward arrows.
Example: Silicon, in its neutral atom, has 14 electrons. Thus, there 2 electrons in 1s, 2 in 2s, 6 in 2p, and 2 more electrons in 3s.
The letters "A" , "B", etc shows the steps taken to write the electron configuration for silicon. Note that in the 3p subshell, there are two electrons that occupy separate orbitals without being paired. This is because of Hund's Rule - when electrons occupy orbitals of equal energy, they don't pair up until they have to. So one would fill up each orbital with the "up" arrows first before adding "down" arrows.
You would write it as:
1s22s22p6 3s23p2
FOR IONS:
Negatively charged ion:
- starting from the number of electrons in a neutral atom (equal to atomic #), add electrons according to the ion charge
- with the new number of electrons, write the electronic configuration
- remember that if there were half-filled orbitals, fill the remaining down arrows in!
Eg. Oxygen ion - O-2
Number of electrons = 10
↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s 2s [ 2p ]
Therefore: 1s22s22p6
Positively charged ion:
- starting from neutral number of electrons, remove electrons (from the outermost shell first working backwards) according to charge
- work your way backwards with the orbitals, so if you have electrons in both s- and p-orbitals, remove the ones in the p-orbital first
- then write the notation the same way you would for the neutral atom but with a different number of electrons!
Core Notation
- Set of electrons in an atom can be divided into 2 subsets: core and outer electrons. Thus core notation shows the electronic configuration in terms of core and outer electrons
- CORE: set of electrons with the configuration of the first noble gas that comes BEFORE it, or the way I like to think of it, the noble gas that's a row above the element
- OUTER: the remaining electrons outside of the core
How do you write it? Well, you will need to refer to your periodic table!
- Find the element.
- Then find the noble gas on the row above it.
- Replace the first part of the configuration, that equals the same amount of electrons as the noble gas, with the symbol of the noble gas in square brackets.
- Leave the remaining electronic configuration of the outer electrons outside the brackets.
eg. Oxygen ion - 10 electrons
1s22s22p6
On the periodic table, Helium is the first noble gas that comes before it. Helium has 2 electrons, so you would replace 1s^2.
Rewrite the configuration and you would get:
[ He ]2s22p6
The red indicates the core and the blue indicates the outer electrons!
Also there are TWO exceptions you need to memorize!!
- Copper (Cu) gains stability by having a full d-subshell, so therefore its configuration would be [Ar]4s3d10
- Chromium (Cr) gains stability by having a half-full d-subshell so it would be written as [Ar]4s3d5
Here is an additional periodic table with some of the configurations already! Your welcome! :)
| 1A | 2A | 3A | 4A | 5A | 6A | 7A | 8A |
| 1 H 1s1 | 2 He 1s2 | ||||||
| 3 Li 1s2 2s1 | 4 Be 1s2 2s2 | 5 B 1s2 2s22p1 | 6 C 1s2 2s22p2 | 7 N 1s2 2s22p3 | 8 O 1s2 2s22p4 | 9 F 1s2 2s22p5 | 10 Ne 1s2 2s22p6 |
| 11 Na [Ne] 3s1 | 12 Mg [Ne] 3s2 | 13 Al [Ne] 3s23p1 | 14 Si [Ne] 3s23p2 | 15 P [Ne] 3s23p3 | 16 S [Ne] 3s23p4 | 17 Cl [Ne] 3s23p5 | 18 Ar [Ne] 3s23p6 |
Oh my.... looking back this was a VERY LONG post. I'm sorry for those readers that had to endure all this way!
HAHAHAHA. In no way take this seriously! I apologize in advance!
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