Electron Configuration Atomic Orbital – region in space in which the probability of finding an electron is high. n Principle Quantum Number – describes the main energy level an electron occupies. l
Subsidiary Quantum Number (azimuthal) – designates the shape of the s region in space an electron occupies. d
p l = 0, 1, 2, 3, … (n-­‐1) s p d f Increasing Energy Level f ml Magnetic Quantum Number – designates the spacial orientation of an atomic orbital. ml = (-­‐l), …, 0, … , (+l) ms Spin Quantum Number – refers to the spin of an electron and orientation of the magnetic field produced by this spin. For every combination of n, l, ml there is !
an ms = ± ! Aufbau Order n Examples: Notation \ 1s 2s 2p 1s Element C 1s22s22p2 __ __ __ __ __ 2s 2p O [He] 2s22p4 __ __ __ __ __ 3s 3p 3d 0 0 -­‐1 0 1 ml 4s 4p 4d 4f Note: i) The symbol represents the individual electron and its spin associated with ms. 5s 5p 5d 5f ii) The abbreviation [He] is a shortcut to show that 6s 6p 6d all orbitals are completely filled up through the noble gas [He]. 7s 7p iii) No two electrons in an atom may have identical sets of four quantum numbers. The shaded oxygen electron above has the unique set: 2, 1, -­‐1, -­‐1/2 The Pauli Exclusion Principle
The Pauli exclusion principle suggests that only two electrons with opposite spin can
occupy an atomic orbital. Stated another way, no two electrons have the same 4
quantum numbers n, l, m, s. Pauli's exclusion principle can be stated in some other
ways, but the idea is that energy states have limit room to accommodate electrons. A state
accepts two electrons of different spins.
In applying this rule, You should realize that an atomic orbital is an energy state.
Hund's RULE
Hund's rule suggests that electrons prefer parallel spins in separate orbitals of
subshells. This rule guides us in assigning electrons to different states in each sub-shell
of the atmic orbitals. In other words, electrons fill each and all orbitals in the subshell
before they pair up with opposite spins.
Pauli exclusion principle and Hund's rule guide us in the aufbau process, which is
figuring out the electron configurations for all elements.
Special Electronic Configurations When two electrons occupy the same orbital, they
not only have different spins (Pauli exclusion principle), the pairing raises the energy
slightly. On the other hand, a half filled subshell and a full filled subshell lower the
energy, gaining some stability. Bearing this in mind, you will be able to understand why
we have the following special electronic configurations.
Cr [Ar]4s1 3d5 <=All s and d subshells are half full
Cu [Ar]4s1 3d10 <=Prefers a filled d subshell, leaving s with 1
Nb [Kr]5s1 4d4 <=5s and 4d energy levels are close
Mo [Kr]5s1 4d5 similar to Cr above
Tc [Kr]5s2 4d5 (not special, but think of Hund's rule)
Ru [Kr]5s1 4d7 <= Only 1 5s electron
Rh [Kr]5s1 4d8 <= in both
Pd [Kr]5s0 4d10 <= Note filled 4d and empty 5s
Ag [Kr]5s1 4d10 <= partial filled 5s, but filled d
Answers : 1) [Ar] ↑↓ ↑↓ ↑ ↑ ↑ ↑ 2) [Ar] 4s23d64p1 [Ar] 4s23d6 3) 3, 2, -­‐2, ½ 4) Ar Increasing Energy Level 1) Diagram the electron configuration for iron Fe in the chart below and write down its configuration. 2) What is the electron configuration of Ge+1 ? 3) What is the set of four quantum numbers that represent the 3d1 electron? 4) The electron configuration of P-­‐3 is identical to what other elements configuration? [Rn] __ 7s (86) __ __ __ 6p __ __ __ __ __ 5d [Xe] __ 6s __ __ __ __ __ __ __ 4f (54) __ __ __ 5p [Kr] __ 5s __ __ __ __ __ 4d (36) __ __ __ 4p [Ar] __ 4s __ __ __ __ __ 3d (18) __ __ __ 3p [Ne] __ 3s (10) __ __ __ 2p [He] __ 2s (2) __ 1s 1 2 3 l=
0 Orbital Geometry n l
0 ml 1 2 3 {0} 1s 0, 1 {0} 2s {-­‐ 1, 0, 1} 2px 0, 1, 2 2pz 2py {0} 3s {-­‐1, 0, 1} 3px 3dxy 3pz 3py (-­‐2, -­‐1, 0, 1, 2} 3dxz 3dz2 3dyz Created by staff FSCJ South Campus Library Learning Commons, Science Lab 07-­‐19-­‐11. Permission to copy and use is granted to all FSCJ staff provided this copyright label is displayed 3dx2 -­‐ y2