Chap 3 Electron Configurations & Quantum Numbers

Quantum Numbers Help us locate all the electrons.

The number and relative energies of all hydrogen electron orbitals through n=3 At ordinary temperatures essentially all hydrogen atoms are in their ground states The electron may be promoted to an excited state by the absorbtion of a photon with the appropriate quantum of energy

Emission lines As excited electrons relax to their ground state they give off light waves at very specific wavelengths called emission lines

Quantum Mechanical Model Proposed by Schrodinger to account for matters’ wave-like behavior. Estimates Probability of finding an electron in an area 90% of the time. Replaces Bohr’s planetary orbits with orbitals, shown as fuzzy clouds

Quantum Numbers: n, l, m, s n: the primary energy level (quanta) –average distance from nucleus l: the subLevel –s, p, d, f m: the number of orbitals within a sublevel –1,3,5,7 s: the electron Spin –up  & down 

Sublevels Number of sublevels increase as radius increases (as n increases) energy # sublevels name of level n= n sublevels n=11 sublevels n=22 sublevels, p n=33 sublevels, p, d n=44 sublevels, p, d, f

Orbitals The different sublevels can hold different # of Orbitals Sublevel# of Orbitals s1 p 3 d 5 f 7

Orbitals Have specific shapes and quantities

Orbital Shapes

The Electron Configuration Notation

Electron Filling Rules The Aufbau Principle Electrons are added one at a time to the lowest orbital available until all of the electrons are used. The Pauli Exclusion Principle An orbital can have a maximum of two electrons. To occupy the same orbital, two electrons must spin in opposite directions Hund’s Rule Electrons occupy equal energy orbitals so that the maximum number of unpaired electrons result.

Electron Configurations Element# of Electrons in ElementElectron Configuration He21s 2 Li31s 2 2s 1 Be41s 2 2s 2 O81s 2 2s 2 2p 4 Cl171s 2 2s 2 2p 6 3s 2 3p 5 K191s 2 2s 2 2p 6 3s 2 3p 6 4s 1

Chlorine Electron Configuration The electron configuration for chlorine is 1s 2 2s 2 2p 6 3s 2 3p 5 The large numbers represent the energy level. The letters represent the sublevel. The superscripts indicate the number of electrons in the sublevel.

Filling using an Aufbau Diagram

1s11s1 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87

He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86 1s21s2 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6

Electron Configurations repeat The shape of the periodic table is a representation of this repetition. When we get to the end of the column the outermost energy level is full. This is the basis for our shorthand.

Yes there is a shorthand!

The Shorthand Write the symbol of the noble gas before the element. Then the rest of the electrons. Aluminum - full configuration. 1s 2 2s 2 2p 6 3s 2 3p 1 Ne is 1s 2 2s 2 2p 6 so Al is [Ne] 3s 2 3p 1

More examples Ge = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 Ge = [Ar] 4s 2 3d 10 4p 2 Hf=1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 2 Hf=[Xe]6s 2 4f 14 5d 2

Writing Electron configurations Using the periodic table

Alkali metals all end in s 1 Alkaline earth metals all end in s 2 really have to include He but it fits better later. He has the properties of the noble gases. s2s2 s1s1 S- block

The P- block p1p1 p2p2 p3p3 p4p4 p5p5 p6p6

Transition Metals -d block d1d1 d2d2 d3d3 d 4 d5d5 d6d6 d7d7 d8d8 d 9 d 10

F - block inner transition elements

Each row (or period) is the energy level for s and p orbitals

D orbitals fill up after previous energy level so first d is 3d even though it’s in row d

f orbitals start filling at 4f f 5f

The Shorthand Again Sn- 50 electrons The noble gas before it is Kr [ Kr ] Takes care of 36 Next 5s 2 5s 2 Then 4d 10 4d 10 Finally 5p 2 5p 2

Configuration of Ions Ions always have noble gas configuration. Na is 1s 1 2s 2 2p 6 3s 1 Forms a +1 ion - 1s 1 2s 2 2p 6 Same configuration as neon. Metals form ions with the configuration of the noble gas before them - they lose electrons.

Configuration of Ions Non-metals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them.

The electrons in the outermost shell (the ones with the highest value of n) are the most energetic, and are the ones which are exposed to other atoms. This shell is known as the valence shell. The inner, core electrons (inner shell) do not usually play a role in chemical bonding. Elements with similar properties generally have similar outer shell configurations. For instance, we already know that the alkali metals (Group I) always form ions with a +1 charge; the "extra" s 1 electron is the one that's lost: IALi1s 2 2s 1 Li+1s 2 Na1s 2 2s 2 2p 6 3s 1 Na+1s 2 2s 2 2p 6 K1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 K+1s 2 2s 2 2p 6 3s 2 3p 6

The Group IIA and IIIA metals also tend to lose all of their valence electrons to form cations. IIABe1s 2 2s 2 Be 2+ 1s 2 Mg1s 2 2s 2 2p 6 3s 2 Mg 2+ 1s 2 2s 2 2p 6 IIIAAl1s 2 2s 2 2p 6 3s 2 3p 1 Al3+1s 2 2s 2 2p 6 The Group IV - VII non-metals gain electrons until their valence shells are full (8 electrons). IVAC1s 2 2s 2 2p 2 C4-1s 2 2s 2 2p 6 VAN1s 2 2s 2 2p 3 N3-1s 2 2s 2 2p 6 VIAO1s 2 2s 2 2p 4 O2-1s 2 2s 2 2p 6 VIIAF1s 2 2s 2 2p 5 F-1s 2 2s 2 2p 6

The Group VIII noble gases already possess a full outer shell, so they have no tendency to form ions. VIIIANe1s 2 2s 2 2p 6 Ar1s 2 2s 2 2p 6 3s 2 3p 6

Table of Allowed Quantum Numbers nlmlml Number of orbitals Orbital Name Number of electrons 10011s1s s2s2 1-1, 0, +132p2p s3s2 1 33p3p6 2-2, -1, 0, +1, +253d3d s4s2 1-1, 0, +134p4p6 2-2, -1, 0, +1, +254d4d10 3-3, -2, -1, 0, +1, +2, +374f4f14