Unit 2 Electrons and Periodic Trends
AP Chemistry
A. Quantum Mechanical Model- Bohr’s model of the
atom worked well for hydrogen, but it seemed to fail with
larger atoms. It was discovered that the “wave” nature of
electrons better explains structure of atom.
1.
Louis de Broglie in the mid 1920’s proposed that
electrons behave with wave and particle properties at
the same time.
a. de Broglie’s equation predicts all moving particles
have wave characteristics:
𝒉
𝝀 = 𝒎∙𝒗
2.
Werner Heisenberg- it is impossible to know both the
position and velocity of an electron simultaneously.
a. velocity of an electron is related to it’s wave nature,
think deBroglie’s equation.
b. position is related to it’s particle nature, think about
Bohr’s model.
c. an electron is observed to be either a particle or a
wave, but never both at once.
3.
Erwin Schrodinger refined the wave-particle theory
proposed by de Broglie.
a. Developed an equation (the wave function) that
treated an electron like a wave and predicted the
probable location of an electron around the
nucleus called an atomic orbital.
b. The atomic model in which electrons are treated
like waves is called the quantum mechanical
model.
4.
Quantum mechanics describes the probable location of
electrons in atoms: energy level, sublevel, orbital, spin.
5.
The overall size and energy of an electron is described
by the principal energy levels (n), sometimes called
shells. We label them with integers: n=1,2,3,4 etc.
Determines the overall size and energy of an orbital.
a. Larger the integer=greater distance from nucleus.
b. Greater distance=less tightly bound.
−𝟐.𝟏𝟖𝒙𝟏𝟎−𝟏𝟖 𝑱
c.
𝑬𝒏 =
d.
Use 2n2 to determine # of e-.
𝒏𝟐
6.
Energy levels are broken down into Sublevels,
sometimes called subshells, define the orbital shape.
The first energy level has one sublevel, the second
energy level has two sublevels, the third three etc.
7.
Atomic orbitals- defines 3-D spatial orientation. They
are labeled s, p, d, and f. The letters refer to shapes
only.
a. number of orbitals: s (1), p (3), d (5), f (7)
8.
Electron spin- property of an electron that causes it to
behave as though it was spinning on an axis.
a. thereby generating a magnetic field,
b. whose direction depends on the direction of the
spin. No two electrons in same orbital can have the
same spin.
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B. Electron Arrangements in Atoms- also called the
shell model or just electron configuration. Three rules
determine an electrons arrangement:
1.
2.
3.
4.
5.
Aufbau Principle- each electron occupies the lowest
energy orbital available.
Pauli exclusion principle- a maximum of two electrons
may occupy a single atomic orbital, but only if the
electrons have opposite spin.
Hund’s rule- a single electron with the same spin can
occupy each orbital then they can pair up with an
opposite spin electron.
Electron configuration- listed in order of filling.
a. principal energy level (n=1, n=2, etc.) is written first
b. atomic orbital is written next; s, p, d, f
c. number of electrons in that orbital is written using a
superscript
d. abbreviated: replace inner (non-valence) electrons
with noble gas symbol, e.g. Al:[Ne]3s23p1
Rearrangement of electrons in order to enhance stability.
a. Electron filling is out of order for groups 6 and 11
where an s electron moves to the d sublevel in
order to half fill or completely fill the d sublevel.
b. Half or completely full sublevels are more stable and
therefore at a lower energy state.
c. Cr:[Ar]4s13d5 or Cu:[Ar] 4s13d10
6.
Valence electrons- electrons that are in an atom’s
outermost energy levels that determine an elements
chemical properties. Associated with an atom’s highest
principle energy level and usually s or p orbitals.

These electrons involved with forming chemical
bonds.
7.
Electron-dot diagram- an element’s symbol, which
represents the nucleus and inner-level electrons (core
electrons), surrounded by dots representing the atom’s
valence electrons.
8.
PES- Photo Electron Spectroscopy. Analytical technique
that provides direct evidence for the shell model.
a. determines energy needed to eject electrons from
atoms, that allows one to infer the e- configuration.
b. x-axis shows binding energy MJ/mole (energy
needed to remove the e-).
c. the greater the binding energy the stronger the
attraction from the nucleus and the closer the e- are
to nucleus.
d. y-axis shows relative # of e-.
C. Organization of the periodic table1.
The PT is organized into four blocks corresponding to
the filling of the four quantum sublevels: s, p, d, f
a. Row (period)- equals the highest principal quantum
number. Also same as the valence energy level.
b. Column (group)- same #of valence electrons.

the number of groups in a block corresponds to max
# of e- that can occupy that sublevel.
c. group 1 → alkaline metals
group 2 → alkaline earth metals
group 18 → noble gases
group 17 → halogens
groups 3-12 → transition metals
lanthanide and actinide series → inner transition metals
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2.
Metals, nonmetals, and metalloids.
a. metals left side of stair step.

shiny, conduct heat and electricity, malleable and

ductile, mostly solids (except Hg)

form ionic compounds with nonmetals

small positive ionization energy

positive or small negative electron affinity

lose electrons during reactions

(alkali metals are most reactive)
b.
c.
nonmetals- to the right of the stair step and H.

opposite properties of metals

form molecules in addition to ionic compounds.

large positive ionization energy

large negative electron affinity, except groups
15 and 18

gain or share electrons during reactions, except
noble gases. Halogens are most reactive.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
1
2
3
4
5
6
7
?
lanthanide
actinide
Metals
Metalloids
?
Nonmetals
metalloids- touch the stair step, except Al
intermediate properties depending on physical and
chemical conditions
D. Periodic Trends- chemical properties of elements are
determined by the # of valence electrons. Properties are
periodic because the number of valence electrons is periodic.
1.
Atomic Radius- The electron cloud surrounding the
nucleus is based on probability and does not have a
clearly defined edge.
a. Radius is defined as half the distance between
nuclei of identical atoms that are bonded.
b. Radius determined by the strength of attraction
between the valence electrons and the nucleus.
c. Force attracting e- and p+ is as a result of
Coulombs Law:
greater the charge= stronger the attraction
greater the distance= weaker the attraction
𝐸=𝑘
2.
𝑞1∙ 𝑞2
𝑑

effective nuclear charge (Zeff) is the charge felt
by the valence electrons after you have taken
into account the number of shielding electrons
that surround the nucleus.

Zeff ∝ (# p – # core e-)

The electron shielding effect is the effect
where core electrons block valence electrons
from the nuclear charge of the nucleus.
d.
Moving down a group, # of energy levels increases,
thus size of electron cloud increases. Atomic radius
increases.
e.
Moving across a period the outer energy level
remains the same, yet number of protons increases
(Zeff increases), as a result the electron cloud is
pulled in tighter. Atomic radius decreases.
Ionic Radius- Radius of an ion
a. When a neutral atom gains e- to become a negative
ion (anion), radius increases; #protons remain
same, yet electron cloud increases in size.
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3.
b.
Neutral atom converted to a positive ion (cation),
losing an electron causes electron cloud to be
pulled tighter by existing protons.
c.
Cation is smaller compared to it’s neutral atom;
opposite true for anion.
Ionization Energy- minimum energy required to remove
an electron from a gaseous atom to form an ion. (kj/mol)
X(g)
X+(g) + ea. always a + value, greater value=harder to ionize)
b. Inversely proportional to atomic radius
c. Increases across a period. # of protons increasing,
therefore; Zeff increases.
d. Decreases down a group, due to shielding effect of
inner electrons.
e. Anomalies: concept= filled and ½ filled sublevels
are particularly stable, requires more energy to
remove.
1. group 13: ionized e- comes from p vs. s
(1s22s22p1)
2. group 16: ionized e- comes from a full orbital
(higher energy than ½ filled) 1s22s22p4
f.
Successive ionization energies:
*small increases within a sublevel
**greater increase between sublevels
***greatest increase between energy levels
4.
Electron Affinity- the energy given off when a neutral
atom in the gas phase gains an extra electron to form a
negatively charged ion.
a. Negative value means energy released, atom more
stable.
b. Positive value means energy is absorbed and the
atom is at a higher energy state, which is unstable
and unlikely to form.
c. The electron affinity is a measure of the attraction
between the incoming electron and the nucleus - the
stronger the attraction, the more energy is released.
d. Trend: increases across a period. Filled and ½ filled
sublevels have an effect.
5.
Electronegativity- indicates the relative ability of an
atom to attract electrons in a chemical bond.
1. Determines the type of bonding between
atoms.
2. Generally decreases as you move down a
group, and increases as you move left-to-right
across a period.
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