Energy Ordering of Orbitals in a Multi Electron Atom

1. Electron ConfigurationThe way in which electrons are distributed among the various orbitals of an atom.Have a copy of the periodic table in front of you while watching this video

2. Energy Ordering of Orbitals in a Multi Electron Atom • Aufbau Principle: • Aufbau in German means “building up” • Electrons fill in the increasing order of energy. • Electrons are added to an orbital diagram one at a time to the lowest energy orbital available, and that the orbital diagram is thus "built up”. • 1s 2s  2p  3s…..

3. Hund’s Rule • Thelowest energy electron configuration for an atom is one having the maximum number of electrons with parallel spins in degenerate orbitals. • When filling sublevels other than s, one electron goes into each orbital until all the orbitals of the subshell are half-filled (all with usually spin up). Once all the orbitals of the subshell are half-filled the pairing of electrons can take place. p subshell p subshell Correct Incorrect (violation of Hund’s 3 unpaired electrons rule) 1 unpaired electron

4. Pauli Exclusion Principle • No two electrons in the same atom can have identical sets of four quantum numbers n, I, ml, and ms. • The first three quantum numbers (n, l, and ml) specify a particular orbital, such as 1s (1,0,0). The fourth quantum number (ms) specifies the spin of the electron. Since there are only two possible values for ms(+1/2 and −1/2), only two electrons can occupy any given orbital and they must have opposite spins. • When we draw a pair of electrons in an orbital using up and down arrows, if one arrow is up and the other must be down. 1s 1s Correct Incorrect (violation of Pauli Exclusion Principle)

5. Writing Electron Configurations Using the PT • C • S • V • I

6. Orbital Diagrams (Using Box and Arrow) • We often represent an orbital as a square and the electrons in that orbital as arrows • The direction of the arrow represents the relative spin of the electron (up spin as +1/2 and down spin as -1/2) • Orbital diagrams are filled using the Aufbau principle, the Pauli exclusion principle, and Hund's rule. • E.g. The electron configuration of Se is 1s2 2s2 2p6 3s2 3p64s23d104p4 1s 2s 2p 3s 3p 4s 3d 4p Orbital Diagram of Se

7. Writing Condensed (Shorthand) Electron Configurations [Core]+ higher energy electrons Core or Kernel = inner electrons 1) Find the symbol for the element on a periodic table. Example: iron (n =4) 2) Write the symbol in brackets for the noble gas located at the far right of the preceding horizontal row (n = 3) on the table. [Ar] 3) Move back down a row (to the row containing the element you wish to describe). Following the elements in the row from left to right, write the outer-electron configuration. [Ar] 4s2 3d6

8. Shorthand notation practice Examples aluminum: 1s22s22p63s23p1 [Ne]3s23p1 nickel: 1s22s22p63s23p64s23d8 [Ar]4s23d8 iodine: [Kr]5s24d105p5 lead (Pb): [Xe]6s24f145d106p2 In the exams you are expected to write the expanded form.

9. Some Common Exceptions The most common exceptions are d4 and d9. A d subshell that is half-filled or full (i.e. 5 or 10 electrons) is more stable than the s subshell of the next shell. ns2 (n-1)d4 ns1(n-1)d5 (less stable) (more stable) E.g. The electron configuration of Cr and Cu. Cr: 1s22s22p63s23p64s23d4 (incorrect) :1s22s22p63s23p64s13d5 (correct) Cu: 1s22s22p63s23p64s23d9 (incorrect) :1s22s22p63s23p64s13d10 (correct)

10. Practice: Writing Electron Configurations and Orbital Diagrams of atoms 1) Write the electron configuration of Na. • Write the electron configuration of Mo. • Which of the electron configurations is not possible? a) 1s22s2 2p3 b) 1s2 2s22p6 c) 1s2 2s22p2 d) 1s22s32p2 e) 1s2 2s2 2p63s1 4) Write the orbital diagram for O.

11. Practice: Writing Electron Configurations and Orbital Diagrams of atoms • There are _______ unpaired electrons in fluorine. • Write the electron configuration of the following atoms. • Ag b) Ca

12. Outer Shell (Valence Shell) Electron Configurations • Elements in any given group in the periodic table have the same type of electron arrangements in their outermost shells. • E.g. the electron configurations of the halogens fluorine and chlorine are [He]2s22p5 and [Ne]3s23p5, respectively. • Outer shell electron configuration for halogens is ns2np5 • E.g. the electron configurations of the alkaline earth metals magnesium and calcium are [Ne]3s2and [Ar]4s2, respectively. • Outer shell electron configuration for alkali earth metals is ns2

13. Valence Electrons • The electron(s) in the outermost shell of an atom. E.g. Na 1s2 2s2 2p6 3s1 • The inner electrons are called as the “Core electrons” Na 1s2 2s2 2p6 3s1 • Valence electron of atom = Group number of that atom Rules for defining valence electrons. • Electrons in the outer most energy level • For main group (representative) elements (elements in s and p orbitals) electrons in filled d or f shells are not valence electrons Core Electrons Valence electron

14. Valence Electrons Examples: (valence electrons in blue) • Na= [Ne]3s1 I: [Kr]5s24d105p5 • Sr = [Kr]3s2 N: [He]2s22p3 • As: [Ar] 4s23d104p3 Practice: 1. What is the outer shell electron configuration of • Alkali earth metals? __________________ • Noble gases? ________________ 2. What is the valence-shell electron configuration for the third period element in Group IIIA?

15. Broad Periodic Table Classifications • Group 1 – Alkali metals (ns1) • Group 2 – Alkaline earth metals (ns2) • Group 7 – Halogens (ns2np5) • Group 8 – Noble gases or Inert gases (ns2np6) • d block elements are transition metals. • 4f block elements are lanthanides. • 5f block elements are called actinides.

16. Practice: Periodic Table 1. Mn, manganese is a/an a) alkali metal b) transition metal c) nonmetal 2. Aluminum is a/an a) alkali metal b) transition metal c) nonmetal d) metal e) none of the above 3. Group 2A elements possess a) 1 outer-shell s electron b) 2 outer-shell s electron c) 2 outer-shell p electrons d) 1 outer-shell d electron 4. The actinides owe their position in the periodic table to the filling of _____ orbitals. (5d/6d/4f/5f)