Transcript Electron Configuration

Electron Configuration
 Energy Levels (1-7): looking at the periodic table you can tell
how many energy levels an atom has by looking at the row the
electron is in.
 Sublevels (s,p,d,f): looking at the periodic table you can tell
sublevels based on the 4 blocks the periodic table is broken up
into.
 Orbitals (3-D orientations):
 s_
 p___
 d_____
 f_______
Rules to know by name….
 Hund’s Rule:
 When several orbitals of equal energy are available, as in a
given sublevel, electrons enter singly with parallel spins
(up arrows first).
 Aufbau Principle:
 The principle postulates a hypothetical process in which an
atom is "built up" by progressively adding electrons. As they
are added, they assume their most stable conditions with
respect to the nucleus and those electrons already there
(lowest energy level first).
 When writing electron configurations:
 We use electrons to represent arrows.
 Each orbital can only hold a maximum of 2 electrons.
Quantum
Numbers
 The principal quantum number:
 Symbolized by n, basically the energy level the electron is in.
 The orbital quantum number:
 Symbolized by l, basically represents the sublevel the electron is in s,p,d, or f.
values for l: s = 0, p: l = 1, d: l = 2, f: l = 3
*The letters s,p,d, and f come from the adjectives used to describe spectral lines: sharp, principal,
diffuse, fundamental.
 The magnetic quantum number:
 Symbolized by m l, this determines the direction in space of the electron cloud
surrounding the nucleus.
 All of the orbitals in a sublevel have the same energy
s __
p __ __ __
d __ __ __ __ __ f __ __ __ __ __ __ __
 The spin quantum number:
 Symbolized by ms, this represents the electron spin.
 ms can equal +1/2 (up arrow) or -1/2 (down arrow)
Examples
 Give the quantum numbers for the outermost electron
in neon, copper, and barium.