Periodic Table Trends & Electron Configuration
Electron Configuration
• A method used by scientists to show the distribution of e in
their orbital.
• Diagonal Rule
2s2ámax # e it holds
Energy levelÆ2s2
2s2
ã
shape of orbital
• Energy levels: the region around the nucleus where the e is likely to be moving.
There are 7 levels.
• Sublevels: represented by the letters s,p,d,f
• Each can hold a maximum
# of electrons
sublevel #electrons
s 2
p 6
d 10
f 14
• Constructing
1. Determine the #e in the element.
2. Begin w/ 1s & follow the arrows for the order of filling energy
levels w/ e
3. The sublevel can only hold a certain # electrons
4. Keep a tally of distributed electrons
Short Hand Electron Configuration
1. Determine the previous noble gas of the element & place its symbol in brackets
2. Determine the difference in the # electrons between the element and the noble gas.
3. Begin with the sublevel “s” in the energy level of the element.
4. Distribute the remaining e according to the diagonal rule.
• Tricks:
Group 1: end in s1
Group 2: end in s2
Group 3: end in p1
Group 4: end in p2
Group 5: end in p3
Group 6: end in p4
Group 7: end in p5
Group 8: end in p6
The number in front of the sublevel is the row number.
Transition metals are the d sublevel.
Movement of Electrons
1. Electrons in a particular path have a fixed energy
2. Electrons change energy by going from one energy level to another when energy (photons) are abs. or emitted
3. Photon abs. energy electron moves away from the nucleus
Low àHigh
4. Photons are emitted electrons move closer to the nucleus
High à Low
• Energy Level: The region around the nucleus where the electron is likely to be moving
• Ground Level: Atoms (electrons) in the lowest energy state
• Quantum Energy: The amount of energy required to move an electron from its present energy level to the next higher one.
Theory Behind Electron Configuration
• Quantum Numbers: The address of an electron to indicate the probable location of the electron within the atom
• n: Principle Quantum #
1. Describes the energy level within the atom.
• Energy levels are 1 to 7
• l:Sublevel
1. Value 0 to n-1
• m: Describes the orbital within a sublevel
1. Which orbital the electron is in
values –l to + l
• ms:Spin
1. Describes the spin of the electron.
• 2. Electrons in the same orbital must have opposite spins.
– Possible spins are clockwise (é) or counterclockwise (ê).
• Pauli’s Exclusion Theory: States no two electrons in an atom can have the same set of four quantum numbers
• Aufau Principle:
States an electron occupies the lowest energy orbital first
• Hunds Rule: The most stable arrangement of electrons in sublevels is one with the greatest # of parallel spins.
Rules Behind the Theory
• C: 1s22s22p2
áâ áâ á á __
1s2 2s2 2p2
NOT
áâ áâ áâ __ __
1s2 2s2 2p2
Electronegativity
• Trend
-Increases across a
period
-Decreases down a
group
• Fluorine (F) is the most
electronegative
• Francium (Fr) is the
least electronegative
• Reasons for Trend:
1. Increase across a
period b/c of poor
shielding & increasing
protons in the nucleus.
Grabs e elsewhere
• a. Shielding: inner
orbitals “shield” e on higher levels from the pull of the nucleus.
• 2. Decreases down a
group b/c good shielding
& additional e become
more difficult to hold
on to.
Atomic Radius
• One-half the distance between the nuclei of identical atoms that are bonded together
• Trend
- Decrease across
a period
-Increase down a
Group
• Reasons for Trend:
1. Decrease across a
period b/c protons
in the nucleus pull e
more tightly towards
them.
• 2. Increase down a
group b/c energy
levels increase the
distance of the e
from the nucleus- e
don't feel added pull
of the nucleus.
• 3. Shielding blocks the
attraction of the
nucleus for outer e
\ size increases
down a group.
• Atoms & Ions
1. Cation: Smaller than parent; remove e
2. Anion: Larger than
parent; add e
Ionization Energy
• Trend
-Increase across a
period
-Decreases down a
group
• Reasons for trend:
1. e difficult to remove
across a period b/c
shielding is poor &
protons exert a
strong attraction.
Ionization Energy
• 2. e easier to remove
down a group b/c
of good shielding each
time you are one
energy level further
from the nucleus.
Types of Bonds
• 1. Ionic Bonds
a. transfer of e;
involves formation
of ions
cation (+) and
anion (-)
b. Occurs between
a metal & non-
metal
c. Only valence e
(outer most e) are
involved in bonding
d. Valence e are shown
in e dot structures;
valence e = group #
• 2. Covalent Bonds
a. sharing e btwn
two non-metals
b. Two Types
1. Nonpolar:
2 atoms
share e equally
(occurs only if the
2 atoms are
identical) ie: N2, O2
2. Polar:
Occurs when
there is an unequal
sharing of e in the
bond b/c of the
difference in
electronegativity.