Unit 2 Electrons and Periodic Trends
Honors Chemistry
Atomic spectroscopy and the Bohr model- (5.1, 5.2)
I.
A new model of the atom evolved out of the similarities
discovered between the behavior of light & electrons. Analysis of
the light revealed that an elements chemical behavior is related
to the arrangement of it’s electrons.
A.
Wave Nature of Light. Light is a form of electromagnetic
radiation with three characteristics:
1. wavelength- measured in meters or nanometers (m or
nm) is the distance between two consecutive crests.
2. frequency – measured in hertz (Hz) is the number of
wavelengths that pass a certain point per second.
3. speed- how fast a wave is moving through space. All EM
radiation travels at 3.0 x 108 m/s.
4. Because light moves at a constant speed there is a
relationship between frequency and wavelength.
𝒄=𝝀∙𝝂
5.
𝜈 = frequency in Hertz
c = speed of light 3.0 x 108m/s
λ = wavelength in meters
Light energy comes in packets, called photons.
𝑬𝒑𝒉𝒐𝒕𝒐𝒏 = 𝒉 ∙ 𝝂
𝒉=
𝑷𝒍𝒂𝒏𝒄𝒌′ 𝒔
𝒄𝒐𝒏𝒔𝒕𝒂𝒏𝒕 = 𝟔. 𝟔𝟐𝟔 × 𝟏𝟎−𝟑𝟒 𝑱 ∙ 𝒔
substitute �=�∙�
𝑬𝒑𝒉𝒐𝒕𝒐𝒏 =
𝒉∙𝒄
𝝀
B.
By passing light through a prism, the color components of the
light can be separated.
1. A continuous spectrum shows all the wavelengths of
light that are being emitted by white light. (think of a
rainbow)
2. An emission spectrum shows the specific frequencies of
light emitted by a specific atom that is being excited.
3. Atoms can be identified by the light they emit, by their
unique emission spectrum.
C.
The Danish scientist Niels Bohr (1885-1962) explained the
formation of emission spectra (for hydrogen only):
1. Potential energy of an electron depends on its distance
from the nucleus.
2. When an atom absorbs a photon of light, it is absorbing
energy.
a. Absorption of a photon causes a low potential energy
electron in an atom to become a high potential energy
electron.
b. When a high potential energy electron loses some of
its energy, the electron moves closer to the nucleus
and the energy lost is emitted as a photon.
3.
Since light energy is quantized, the energy of an electron
must also be quantized. In other words, an electron
cannot have just any amount of potential energy.
a. Within the atom there must be a number of distinct
energy levels, analogous to steps on a staircase.
b. Where you are at on the “staircase” is restricted to
where the stairs are. Similarly, there are only a
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limited number of permitted energy levels in an atom.
An electron cannot exist between levels.
4.
Equation to calculate the energy that an electron would
have at any energy level:
−𝟐. 𝟏𝟖𝒙𝟏𝟎−𝟏𝟖 𝑱
𝑬𝒏 =
𝒏𝟐
a.
b.
c.
n is the energy level in question and the negative
sign means that the lower energies correspond to
states with larger negative numbers for energy
values – be careful!
ground state (n = 1) electron has lowest (most
negative) energy
excited state (n > 1), electron energy increases until
ionized (E = 0 J)
∆Eelectron = En-final – En-initial
∆Eelectron > 0 when increasing n
∆Eelectron < 0 when decreasing n
|∆Eelectron| = Ephoton
Bohr developed a conceptual model in which an electron moving
around the nucleus is restricted to certain distances from the nucleus,
these distances are determined by the amount of energy the electron
has. This is called the planetary orbital model.
II. Quantum Mechanical Model- (5.2) 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.
A.
Louis de Broglie in the mid 1920’s proposed that electrons
behave with wave and particle properties at the same time.
1. de Broglie’s equation predicts all moving particles have
𝒉
2.
wave characteristics: 𝝀 =
𝒎∙𝒗
Einstein’s “Photoelectric Effect” suggests electrons have
particle properties as well.
B.
Werner Heisenberg- it is impossible to know both the
position and momentum of an electron simultaneously.
1. velocity of an electron is related to it’s wave nature,
think deBroglie’s equation.
2. position is related to it’s particle nature, think about
Bohr’s model.
3. an electron is observed to be either a particle or a wave,
but never both at once.
C.
Erwin Schrodinger refined the wave-particle theory proposed
by de Broglie.
1. 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.
2. The atomic model in which electrons are treated like
waves is called the quantum mechanical model.
D.
Quantum mechanics describes the probable location of
electrons in atoms: energy level, sublevel, orbital, spin.
E.
The overall size and energy of an electron is described by the
principal energy levels (n), sometimes called shells. We label
2
them with integers: n=1,2,3,4 etc. Determines the overall size and
energy of an orbital.
1. Larger the integer=greater distance from nucleus.
2. Greater distance=less tightly bound.

Use 2n2 to determine # of e- found on any energy level.
F.
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.
G.
Atomic orbitals- defines 3-D spatial orientation. They are
labeled s, p, d, and f. The letters refer to shapes only.
1. number of orbitals: s (1), p (3), d (5), f (7)
H.
Electron spin- property of an electron that causes it to
behave as though it was spinning on an axis; thereby
generating a magnetic field,
1. No two electrons in same orbital can have the same
spin.
III. Electron Arrangements in Atoms- (5.3) also called the shell
model or just electron configuration. Three rules determine an
electrons arrangement:
1.
2.
3.
4.
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 according to
Aufbau principle.
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
5.
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.
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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.
determines energy needed to eject electrons from atoms,
that allows one to infer the e- configuration.
x-axis shows binding energy MJ/mole (energy needed to
remove the e-).
the greater the binding energy the stronger the attraction
from the nucleus and the closer the e- are to nucleus.
y-axis shows relative # of e-.
4.
5.
6.
7.
IV. Organization of the periodic table- (6.1-6.2)
A.
The PT is organized into four blocks corresponding to the
filling of the four quantum sublevels: s, p, d, f
1. Row (period)- equals the highest principal quantum
number. Also same as the valence energy level.
2. Column (group)- same #of valence electrons.

the number of groups in a block corresponds to
max # of e- that can occupy that sublevel.
3. 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
B.
Metals, nonmetals, and metalloids.
1. 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)
2.
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.




3.
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
III. Periodic Trends- (6.3) chemical properties of elements are
determined by the # of valence electrons. Properties are periodic
because the number of valence electrons and their e- configurations
are periodic.
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A.
Atomic Radius- The electron cloud surrounding the nucleus is
based on probability and does not have a clearly defined edge.
1. Radius is defined as half the distance between nuclei of
identical atoms that are bonded.
2. Radius determined by the strength of attraction between
the valence electrons and the nucleus.
3. 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=




B.
C.
𝑞1∙ 𝑞2
𝑑
E is potential energy
q is magnitutde of charge
d is separation distance

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.
4.
Moving down a group, # of energy levels increases, thus
size of electron cloud increases, shielding e- increase.
Atomic radius increases.
5.
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
1. When a neutral atom gains e- to become a negative ion
(anion), radius increases; #protons remain same, yet
electron cloud increases in size.
2.
Neutral atom converted to a positive ion (cation), losing an
electron causes electron cloud to be pulled tighter by
existing protons.
3.
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) + e1. always a + value, greater value=harder to ionize)
2. Inversely proportional to atomic radius
3. Increases across a period. # of protons increasing,
therefore; Zeff increases.
4. Decreases down a group, due to shielding effect of inner
electrons.
5. Anomalies: concept= filled and ½ filled sublevels are
particularly stable, requires more energy to remove.
a. group 13: ionized e- comes from p vs. s (1s22s22p1)
b. group 16: ionized e- comes from a full orbital (higher
energy than ½ filled) 1s22s22p4
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6.
Successive ionization energies:
*small increases within a sublevel
**greater increase between sublevels
***greatest increase between energy levels
D.
Electron Affinity- the energy given off when a neutral atom in
the gas phase gains an extra electron to form a negatively
charged ion.
1. Negative value means energy released, atom more stable.
2. Positive value means energy is absorbed and the atom is
at a higher energy state, which is unstable and unlikely to
form.
3. 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.
4. Trend: increases across a period. Filled and ½ filled
sublevels have an effect.
E.
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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