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UNIT 3
PERIODIC TRENDS and
ELECTRON
CONFIGURATIONS
Periodic Table Trends
Periodic Law
“When arranged by increasing atomic number, the chemical
elements display a regular and repeating pattern of chemical and
physical properties.”
Atoms with similar properties appear in groups or families (vertical
columns) on the periodic table.
all have the same number of valence (outer shell) electrons, which
governs their chemical behavior.
Atomic Radius
Radius is the distance from the center of the nucleus to
the “edge” of the electron cloud.
Since a cloud’s edge is difficult to define, scientists use
covalent radius, or half the distance between the nuclei
of 2 bonded atoms.
Atomic radii are usually measured in picometers (pm) or
angstroms (Å). An angstrom is 1 x 10-10 m.
Ex. Two Br atoms bonded together are 2.86 angstroms
apart. So, the radius of each
atom is 1.43 Å.
2.86 Å
1.43 Å 1.43 Å
The trend for atomic radius
smaller at the top to larger at the bottom.
add an entirely new energy level to the electron cloud, making the
atoms larger with each step.
The trend across a horizontal period is less obvious.
Each step from left to right adds a proton and an electron (and 1 or 2
neutrons) and electrons are added to existing energy levels.
The effect is that the more positive nucleus has a greater pull on the
electron cloud.
The nucleus is more positive and the electron cloud is more negative.
The increased attraction pulls the cloud in, making atoms smaller as we
move from left to right across a period.
Ionization Energy
If an electron is given enough energy to overcome the forces
holding it in the cloud, it can leave the atom completely.
The atom has been “ionized” or charged. The number of protons
and electrons is no longer equal.
The energy required to remove an electron from an atom is
ionization energy. (measured in kilojoules, kJ)
The larger the atom is, the easier its electrons are to remove.
Ionization energy and atomic radius are inversely proportional.
Electron Affinity
Electron affinity is the energy change that occurs when an atom
gains an electron (also measured in kJ). (how well atoms attract
electrons)
Electron affinity is usually exothermic, but not always.
Electron affinity is exothermic if there is an empty
or partially empty orbital for an electron to
occupy.
If there are no empty spaces, a new orbital or
energy level must be created, making the process
endothermic.
Metals: Metals like to lose valence electrons to form cations to
have a fully stable octet. They absorb energy to lose electrons. The
electron affinity of metals is lower than that of nonmetals.
Nonmetals: Nonmetals like to gain electrons to form anions to have
a fully stable octet. They release energy to gain electrons to form an
anion; thus, electron affinity of nonmetals is higher than that of
metals.
Electronegativity
Electronegativity is a measure of an atom’s attraction for
another atom’s electrons.
It is an arbitrary scale that ranges from 0 to 4.
The units of electronegativity are Paulings.
Generally, metals are electron givers and have low
electronegativities.
Nonmetals are are electron takers and have high
electronegativities.
What about the noble gases?
0
Periodic Trends Rap
Electron Configuration
Quantum Mechanical Model (electron cloud)
Energy is quantized. It comes in chunks.
A quanta is the amount of energy needed to move from one
energy level to another.
Since the energy of an atom is never “in between” there must be
a quantum leap in energy.
Schrödinger derived an equation that described the energy and
position of the electrons in an atom
Atomic Orbitals
Within each energy level the complex math of
Schrödinger's equation describes several
shapes.
These are called atomic orbitals
Regions where there is a high probability of
finding an electron
S orbitals
1 s orbital in
every energy level
1s
Spherical shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals
2s
3s
P orbitals
Start at the second energy
level
3 different directions
3 different shapes
Each orbital can hold 2
electrons
The p Sublevel has 3 p
orbitals
Called the 2p, 3p, etc.
D orbitals
The D sublevel starts in the 3rd energy level
5 different shapes (orbitals)
Each orbital can hold 2 electrons
The D sublevel has 5 D orbitals
Called the 3d, 4d, etc.
F orbitals
The F sublevel starts in the fourth energy level
The F sublevel has seven different shapes
(orbitals)
2 electrons per orbital
The F sublevel has 7 F orbitals
ONLY 4F and 5F
Summary
Sublevel
# of shapes
(orbitals)
Max # of
electrons
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
Starts at energy
level
Atomic Orbitals
The periodic table gives clues to how electrons fill
each energy level (shell)
Principal Quantum Number (n) = the energy level of
the electron. (same as the period it’s in)
Three ways to represent where electrons are
Electron Configuration
Noble Gas Configuration
Orbital Notation
Orbital Filling Rules!!!!
Aufbau principle- electrons enter the lowest energy
first.
This causes difficulties because of the overlap of orbitals
of different energies.
Pauli Exclusion Principle- at most 2 electrons per
orbital - different spins
Hund’s Rule- When electrons occupy orbitals of equal
energy they don’t pair up until they have to .