C.P. Chem Unit 10 Guided Overview
Quantum Mechanics and Electron Configurations
The time around the turn of the 20th century was filled with developments in what we thought the atoms “looked like”. This evolution was caused by the intersection of several major ideas and discoveries. Our perception of the nature of light, the atom and matter itself all changed during this time. Many of the ways you view these things are primarily pre 1900’s thinking, how do we fix that?
History of atomic theory:
DemocritusDaltonJJ ThompsonRutherfordBohr
2 things that rocked the atomic theory world:
1)Line spectra
- Bohr's explanation of line spectra
- ground vs. excited state
2)Quantum theory
- Plank
- Heisenberg
-A mathematician, Schroedinger, provided the next step to us.
-Instead of looking for the electron we use math to determine an area (in 3D a volume) where the electron probably exists.
-We define an area where the electron has a 90% probability of existing. We call this area where the electron can probably be found an orbital. (note the difference between an orbital and Bohr’s orbit)
-What we get for a “picture” of the atom is an electron cloud
-Each orbital has its own characteristic shape
-How electrons are arranged around the nucleus gets a bit complicated now.
-Electrons are first defined by the energy level in which they exist. This is roughly equivalent to Bohr’s energy level and is symbolized by n, (n = 1,2,3,4,5,6,7,…). Energy levels always “fill” from the lowest level up, therefore n = 1 fills, then n = 2 fills etc.
-Energy levels can be divided into sublevels. The number of sublevels per energy level is equal to n. ( 1 sublevel in n = 1, 2 sublevels in n = 2, etc.) Sublevels (l) are labeled with letters (in order from lowest energy up) s,p,d,f,g…
-Sublevels contain orbitals. The number of orbitals per sublevel follows the pattern s has 1 orbital, p has 3, d has 5, f has 7, etc. etc. (remember we get shape from orbitals; p. 365)
-Each orbital can contain up to 2 electrons. These electrons are different from each other in spin. They are given the values of +½ and- ½. They are shown by upward and downward pointing arrows.
-Chart:
Energy Sub- Orbitals (e- added)e- per e- per
Level level sub lev en. lev
Orbital Notation: orbitals are shown as blanks, labeled, with up to 2 electrons
○ ○ ○○○
1s 2s 2p
How electrons “fill” around an atom is governed by 3 basic rules (p. 367,368)
(where do you find the # of e-?)
-spin must be different in the orbital blank (Pauli exclusion principle)
-remember, we fill from lowest to highest energy (Aufbau principle)
-Every orbital in a sublevel is singly occupied before any orbital is doubly occupied(Hund’s rule)
-Energy diagrams:
- Now we can do true orbital notations (just a flat version of the energy diagram)
H○ ○ ○○○ ○ ○○○ ○ ○○○○○ ○○○
1s 2s 2p 3s 3p 4s 3d 4p
He○ ○ ○○○ ○ ○○○ ○ ○○○○○ ○○○
1s 2s 2p 3s 3p 4s 3d 4p
Li○ ○ ○○○ ○ ○○○ ○ ○○○○○ ○○○
1s 2s 2p 3s 3p 4s 3d 4p
C______
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p
Al______
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p
Ca
- How do we know when there is overlap of sublevels?
We can also shorten orbital notations into electron configurations
-What do the superscripts add up to?
Examples:
Ca
Sn
Eu
The Periodic Table
-Our modern periodic table’s origin is credited to the Russian Dmitri Mendeleev
-Mendeleev arranged the elements by order of atomic mass so that elements with like properties were arranged in vertical columns
-Mendeleev left some spaces empty for elements that weren’t discovered but had to be there for his scheme to work, when they were later discovered his work was verified.
-Later work by Mosely showed that these varied by atomic number not mass
-All of these use the idea of periodicity and what was later termed the periodic law:
-The modern periodic table shows us many things:
-Groups or families and their names and numbers
-Metal, nonmetal, and semimetal
-Electron configurations
-Valence electrons
-Note relation to ion formed