Chap 3 Electron Configurations Quantum Numbers Quantum Numbers

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 sub. Level – 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=1 1 sublevel s n=2 2 sublevel s, p n=3 3 sublevel s, p, d n=4 4 sublevel s, p, d, f

Orbitals The different sublevels can hold different # of Orbitals Sublevel s p d f # of Orbitals 1 3 5 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 Element Electron Configuration He 2 1 s 2 Li 3 1 s 22 s 1 Be 4 1 s 22 s 2 O 8 1 s 22 p 4 Cl 17 1 s 22 p 63 s 23 p 5 K 19 1 s 22 p 63 s 23 p 64 s 1

Chlorine Electron Configuration The electron configuration for chlorine is 1 s 2 2 p 6 3 s 2 3 p 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

H Li 1 1 s 1 3 1 s 22 s 1 Na 11 1 s 22 p 63 s 1 K 19 1 s 22 p 63 s 23 p 64 s 1 Rb 37 Cs 55 Fr 87 1 s 22 p 63 s 23 p 64 s 23 d 104 p 65 s 1 1 s 22 p 63 s 23 p 64 s 23 d 104 p 65 s 24 d 10 5 p 66 s 1 1 s 22 p 63 s 23 p 64 s 23 d 104 p 65 s 24 d 105 p 66 s 24 f 145 d 106 p 67 s 1

1 s 2 He 2 1 s 22 p 6 Ne 10 Ar 2 2 6 1 s 2 s 2 p 3 s 3 p 18 1 s 22 p 63 s 23 p 64 s 23 d 104 p 6 Kr 36 1 s 22 p 63 s 23 p 64 s 23 d 104 p 65 s 24 d 105 p 6 Xe 54 1 s 22 p 63 s 23 p 64 s 23 d 104 p 65 s 24 d 10 Rn 5 p 66 s 24 f 145 d 106 p 6 86

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. 2 2 6 2 1 • 1 s 2 s 2 p 3 s 3 p 2 2 6 • Ne is 1 s 2 s 2 p 2 1 • so Al is [Ne] 3 s 3 p

More examples • Ge = 1 s 22 p 63 s 23 p 64 s 23 d 104 p 2 • Ge = [Ar] 4 s 23 d 104 p 2 • Hf=1 s 22 p 63 s 23 p 64 s 23 d 104 p 65 s 2 4 d 105 p 66 s 24 f 145 d 2 • Hf=[Xe]6 s 24 f 145 d 2

Writing Electron configurations Using the periodic table

S-s 1 block s 2 • 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.

The Pblock p 1 p 2 p 3 p 4 p 5 p 6

Transition Metals -d block d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10

F - block • inner transition elements f f 1 f 2 f 3 f 4 f 5 6 f 1 f 7 f 8 f 9 f 10 f 11 f 12 3 f 14

1 2 3 4 5 6 7 • 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 3 d even though it’s in row 4. 1 2 3 4 5 6 7 3 d

1 2 3 4 5 6 7 • f orbitals start filling at 4 f 4 f 5 f

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

Configuration of Ions • Ions always have noble gas configuration. • Na is 1 s 12 s 22 p 63 s 1 • Forms a +1 ion - 1 s 12 s 22 p 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: IA Li 1 s 22 s 1 Li+ 1 s 2 Na 1 s 22 p 63 s 1 Na+ 1 s 22 p 6 K K+ 1 s 22 s 22 p 63 s 23 p 64 s 1

The Group IIA and IIIA metals also tend to lose all of their valence electrons to form cations. IIA Be 1 s 22 s 2 Be 2+ 1 s 2 Mg 1 s 22 p 63 s 2 Mg 2+ 1 s 22 p 6 Al Al 3+ 1 s 22 p 6 IIIA 1 s 22 p 63 s 23 p 1 The Group IV - VII non-metals gain electrons until their valence shells are full (8 electrons). IVA VA VIIA C N O F 1 s 22 p 2 1 s 22 p 3 1 s 22 p 4 1 s 22 p 5 C 4 N 3 O 2 F- 1 s 22 s 22 p 6

The Group VIII noble gases already possess a full outer shell, so they have no tendency to form ions. VIIIA Ne 1 s 22 p 6 Ar 1 s 22 p 63 s 23 p 6

Table of Allowed Quantum Numbers n l ml Number of orbitals Orbital Name Number of electrons 1 0 0 1 1 s 2 2 0 0 1 2 s 2 1 -1, 0, +1 3 2 p 6 0 0 1 3 s 2 1 -1, 0, +1 3 3 p 6 2 -2, -1, 0, +1, +2 5 3 d 10 0 0 1 4 s 2 1 -1, 0, +1 3 4 p 6 2 -2, -1, 0, +1, +2 5 4 d 10 3 -3, -2, -1, 0, +1, +2, +3 7 4 f 14 3 4