Section 11.4
Electron Configurations and Atomic Properties
Objectives
1. To understand how the principal energy levels fill with
electrons in atoms beyond hydrogen
2. To learn about valence electrons and core electrons
3. To learn about the electron configurations of atoms with Z <
18
4. To understand the general trends in properties in the
periodic table
Section 11.4
Electron Configurations and Atomic Properties
A. Electron Arrangements
• Orbital Notation for Carbon
  _ _ __
1s 2s
2p
• Electron Configuration for carbon
1s22s22p2
Section 11.4
Electron Configurations and Atomic Properties
“Orbital Filling Rules”
• Aufbau Principle
Electrons fill the lowest energy orbitals available
• Pauli Exclusion Principle
Orbitals can hold a maximum of two electrons; two
electrons in one orbital must have opposite spins.
• Hund’s Rule
Electrons fill equal energy orbitals one at a time until
each is occupied by one electron; Electrons in
singly occupied orbitals have the same spin.
Section 11.4
Electron Configurations and Atomic Properties
A. Electron Arrangements in the First 18 Atoms on the
Periodic Table
Classifying Electrons
• Valence electrons – electrons in the outermost (highest)
principal energy level of an atom
• Core electrons – inner electrons
• Elements with the same valence electron arrangement show
very similar chemical behavior.
Section 11.4
Electron Configurations and Atomic Properties
B. Electron Configurations and the Periodic Table
• Orbital filling and the periodic table
Section 11.4
Electron Configurations and Atomic Properties
A. Electron Arrangements in the First 18 Atoms on the
Periodic Table
Section 11.4
Electron Configurations and Atomic Properties
B. Electron Configurations and the Periodic Table
• Look at electron configurations for K through Kr
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
 filled outer energy level
Ne
1s22s22p6
Na
Mg
1s22s22p63s1  can lose 1 electron to have
a filled energy level
1s22s22p63s2  loses 2 electrons
F
1s22s22p5
 gains 1 electron
O
1s22s22p4
 gains 2 electrons
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Metals and Nonmetals
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Metals
• Conduct electricity
• Left side of the periodic
table
• Form (+) ions; lose
electrons
Nonmetals
• Do not conduct electricity
• Right side of the periodic
table
• Form (-) ions; gain
electrons
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Atomic Size
•Measured as the atomic
radius (distance from the
center of an atom to its
outermost electrons
Remember:
The locations of electrons
are not exact  atomic
radius is an approximation.
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Atomic Size
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Atomic Radius
Trends:
1. Atoms increase in size as you move down a group on
the periodic table
Why?
• Outermost electrons are in higher energy levels. Orbital
size increases with energy level.
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Structure and the Periodic Table
Trends (cont.) - Atomic Radius
2. Atoms decrease in size as you move from left to
right across the periodic table.
Why?
• It is the result of increasing nuclear charge going
left to right. As the number of protons increases,
there is a stronger attraction for the outermost
electrons. (which are located in the same energy
level. The electrons will be held closer to the
nucleus.
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Ionization Energies
• Ionization Energy – energy required to remove an electron
from an individual atom (gas)
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Ionization Energy
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Ionization Energies
Trends:
1. I.E. increases as you go from left to right on the periodic
table
Why?
• Increasing nuclear charge as you move to the right; the
nucleus has a stronger attraction for the outer electrons
2. I.E increases as you move up a group on the periodic
table.
Why?
• Electrons are closer to the nucleus as you move up a
group; there is a stronger attraction to the nucleus
Section 11.4
Electron Configurations and Atomic Properties
C. Atomic Properties and the Periodic Table
Ionization Energies
I.E. can also be explained in terms of atomic radius.
• The smaller an atom is…the greater the attraction
between the nucleus and outer electrons… and the
greater the I.E.
Section 11.4
Electron Configurations and Atomic Properties
Practice
1. Which element (Cs, Hf, Au) has the smallest
atomic radius?
Au
2. Arrange the following elements in order of
decreasing ionization energy: Li, O, C, K,
Ne, F
Ne, F, O, C, Li, K
3. For each of the following pairs of elements
select the largest atom.
C, O
Sr, Mg
Si, N
Li, Mg