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Chapter 5
Periodicity and Atomic Structure
5.1 Development of Periodic Table

A. Creation of the Periodic table
1.
Ideal example of how scientific theory
comes into being
 2.
Random observation
 3.
Organization of data in ways that make sense
 4.
Consistent hypothesis emerges
 5.
Explains known facts
 6.
Makes predictions about unknown
phenomena

5.1 Periodic Table Con’t
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B. Mendelee's hypothesis about organizing
known chemical information
1.
Met criteria for a good hypothesis
 2.
Listed the known elements by atomic weight
 3.
Grouped them together according to their
chemical reactivity
 4.
was able to predict the properties of
unknown elements – eka – aluminum, eka – silicon

The Early Periodic Table
5.2 Electromagnetic Radiation
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Electromagnetic radiation – forms of radiant energy (light in
all its varied forms)
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Electromagnetic spectrum – a continuous range of
wavelengths and frequencies of all forms of electromagnetic
radiation
5.2 Electromagnetic Radiation
Radiation energy – has wavelike
properties

Frequency (ν, Greek nu) – the number
of peaks (maxima) that pass by a fixed
point per unit time (s-1 or Hz)
Wavelength (λ, Greek lambda) – the
length from one wave maximum to the
next
Amplitude – the height measured from
the middle point between peak and
trough (maximum and minimum)
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Intensity of radiant energy is proportional
to amplitude
Electromagnetic Waves
Speed of Light
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Speed of light (c) – rate of travel of all
electromagnetic radiation in a vacuum
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c = 3.00 x 108 m/s
Wavelength x Frequency = Speed
c = λν
Frequency and wavelength are inversely related
long λ  low ν (nu)
short λ  high v (nu)
Example

Calculate the wavelength, in meters, of radiation
with a frequency of 4.62 x 1014 s-1. What region
of the electromagnetic spectrum is this?
Example
5.3– Electromagnetic Radiation and
Atomic Spectra
Individual atoms give off light
when heated or otherwise
excited energetically

Provides clue to atomic makeup
Consists of only few λ
Line spectrum – series of discrete
lines ( or wavelengths) separated by
blank areas
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E.g. Lyman series in the
ultraviolet region
The energy level of Hydrogen
1.– Particlelike Properties of Electromagnetic Radiation: The Planck Equation
Emission of Energy by Atom

How does atom emit light?
Atoms absorbs energy
 Atoms become excited
 Release energy
 Higher-energy photon –>shorter wavelength
 Lower-energy photon -> longer wavelength

5.4 – Particlelike Properties of Electromagnetic
Radiation: The Planck Equation

Blackbody radiation – the visible glow that
solid objects give off when heated

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Intensity does not continue to rise indefinitely as λ
decreases
– the dependence of the intensity of blackbody
radiation on wavelength at different
temperature
Planck Equation
Planck – energy radiated by a heated object is

quantized

Radiant energy emitted in discrete units or quanta
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The smallest quantity of energy that can be emitted in the form of
electromagnetic radiation
E = hν or
h = 6.626 x 10-34 J•s (Planck’s constant)
unit of E is J/proton
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high energy radiation – higher ν, shorter λ
low energy radiation – lower ν, higher λ
Black Body Radiation
Examples

What is the energy (kJ/mol) of photons of radar
wave with

ν = 3.35 x 108 Hz?
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λ = 2.57 x 102 m
Photoelectric Effect
Photoelectric effect –electrons are ejected from the
surface of certain metals exposed to light of at least a
certain minimum frequency

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irradiating a clean metal surface with light causes electrons to be
ejected from the metal
Einstein – beam of light behaves as if it were compose of
photons (stream of small particles)
Energy of photons: E = hν
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Energy depends only on frequency of photon
Intensity of light beam – measure of number of photons, not
energy
Photoelectric Effect
Light energy can behave as both waves and small particles
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Both matter and energy occur only in discrete units
Atoms – emit light quanta (photons) of a few specific energies
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Give rise to a line spectrum
5.5 – Wavelike Properties of Matter : The
de Broglie Equation
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Einstein – relationship between mass and λ
h
E = mc2
or
E
m =
2
c
=
c
de Broglie – matter can behave in some respects like
light

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Both light and matter are wavelike as well as particlelike.
Relationship between λ of an electron or of any other
particle or object of mass m moving at velocity v
m= h
v
Examples

What is the de Broglie wavelength (in meters) of
a pitched baseball with a mass of 120 g and
speed of 100 mph (44.7 m/s)
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At what speed (in meter per second) must a 145
g baseball be traveling to have a de Broglie
wavelength of 0.500 nm?
5.6 - Quantum Mechanics and the
Heisenberg Uncertainty Principle
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A. Bohr – described the structure
of the hydrogen atom as containing
an electron circling the nucleus
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1.
Specific orbits of the
electrons correspond to specific energy
levels
2.
Quantum number: the
smallest increment of energy
i. Represented the energy
difference of any two adjacent orbits
5.6 - Quantum Mechanics and the
Heisenberg Uncertainty Principle
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B. Schrödinger – quantum mechanical model of atom
1. Abandon idea of an electron as a small particle
moving around the nucleus in a defined path
2. Concentrate on the electron's wavelike properties
C. Heisenberg Uncertainty Principle – both the
position (Δx) and the momentum (Δmv) of an electron
cannot be known beyond a certain level of precision
1.
(Δx) (Δmv) > h
4π
2.
Cannot know both the position and the
momentum of an electron with a high degree of
certainty
5.6 - Quantum Mechanics and the
Heisenberg Uncertainty Principle
3.
If the momentum is known with a high
degree of certainty
i.
Δmv is small
ii.
Δ x (position of the electron) is large
4.
If the exact position of the electron is known
i.
Δmv is large
ii.
Δ x (position of the electron) is small
The Bohr Model of the Atom

the Bohr model created by
Niels Bohr depicts the atom
as a small, positively charged
nucleus surrounded by
electrons that travel in
circular orbits around the
nucleus
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similar in structure to the solar
system, but with electrostatic
forces providing attraction,
rather than gravity
Describe the behavior of
electrons in an atom
5.7 The Wave Mechanical Model of
the Atom
Schröndinger’s quantum mechanical model of atomic
structure is frame in the form of a wave equation;
describe the motion of ordinary waves in fluids.
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i. Wave functions or orbitals (Greek, psi , the
mathematical tool that quantum mechanic uses to
describe any physical system
ii. 2 gives the probability of finding an electron
within a given region in space
5.7 The Wave Mechanical Model
of the Atom
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iii. Contains information about an electron’s
position in 3-D space
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defines a volume of space around the
nucleus where there is a high probability of
finding an electron
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say nothing about the electron’s path or
movement
5.8 The Orbitals
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Orbital: the probability map for hydrogen
electrons
The principal quantum number (n): Shell
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a. describes the size and energy level of the orbital
a.
positive integer (n = 1, 2, 3 …..)
as the value of n increases
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the number of allowed orbital increases
size of the orbital increases
the energy of the electron in the orbital increases
The Orbitals
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As the value of n increases, the number of
allowed orbitals increases and the size of the
orbitals become larger, thus allowing an electron
to be far from the nucleus, because it takes
energy to separate a negative charge from a
positive charge
E.gn = 3  third shell (period #3)
n = 5  fifth shell (period # 5)
n=2
The orbitals
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The angular-momentum quantum number (l)
defines the 3-D shape of orbital
integral value from 0 to n-1
If n = 1, then l =0
n = 2, then l = 0 or 1
n = 3, then l = 0, 1 or 2
and so forth
The orbitals
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orbitals are grouping in group according to the
angular-momentum quantum number l is called
subshells.
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Quantum number l
Subshell notation:
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0
s
1
p
2
d
3
f
4 ….
g
The Orbitals
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. Describes the orbitals
that are occupied by the
electrons in an atom
B. Aufbau principle: in
the ground state, the
electrons occupy the
lowest energy orbital first
Orbitals are grouping in
group according to the
angular-momentum
quantum number l is
called subshells.
Types of orbitals
Notations: s, p, d, f
Subshell s
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S orbital
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Probability of finding an electron
depends only on the distance of
the electron from the nucleus
Differences among s orbital in
different shells
Size increases in higher shells
Electron distribution in outer s
orbital has several different
regions of maximum probability
separated by a node
Node – surface of zero
probability
Intrinsic property of a wave –
zero amplitude at node
Subshell p
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dumbbell shape, are oriented on the 3-principle axes,
x, y, and z
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Electron distribution concentrated in identical lobes
on either side of the nucleus

Nodal plane cuts through nucleus
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probability of finding a p electron near the nucleus is
zero
Subshell p
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have different phases
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crucial for bonding
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Only lobes with same phase can interact to
form covalent bonds

Three p orbitals oriented along the x-, y-, and
z-axes (px, py, pz)
Subshell p
Subshell - d
examples

Indicate whether each of the following
statements about the atomic structure is true of
false.
An s orbital is always spherical in shape
 The 2s orbital is the same size as the 3s orbital
 The number of lobes on a p orbital increases as n
increases. That is, a 3p orbital has mores lobes than
a 2p orbital
 Level 1 has one s orbital, level 2 has two s orbitals,
level 3 has 3s orbitals and so on.
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Summary
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Subslevels (type of orbitals) Present
1s (1)
2s (1)
2p (3)
3s (1) 3p (3)
3d (5)
4s (1) 4p (3) 4d (5) 4f (7)
5.10 The Wave Mechanical Model
Pauli exclusion principle
No two electrons in the same atom can have the
same four quantum numbers (ms)
An orbital can hold at most two electrons
If two electrons are in the same orbital, they must
have the opposite spin
The spin quantum number (ms) refers to the two
possible spin orientation of an electron:
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+1/2 (up-spin) or -1/2 (down-spin)
5.10 The Wave Mechanical Model
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E. Hund’s rule: if more than one orbital with
the same energy level are available, one electron
will occupy each orbital until all are filled, before
putting a 2nd electron in any orbital (section
5.10)
5.12- Electron Arrangements in the First
Eighteen Atoms on the Periodic Table
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Recall: Atomic number (Z)
= # electrons = # protons
Electron configuration:
describes the orbitals that are
occupied by the electrons in
an atom
Orbital diagrams: describe
the orbitals with arrows
representing electrons
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a. Arrows are written up or
down to denote electron’s spin
Example

Write the full electron configuration and orbital
filling diagram for: O, Na, Si, Cr
Electrons Configuration

Shorthand version – give the symbol of the
noble gas in the previous row to indicate
electrons in filled shells, and then specify only
those electrons in unfilled shells
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E.g Shorthand version of P: [Ne] 3s2 3p3
The valence-shell electrons are the outer
most shell of electron
E.g
Valence electrons of P is 5
5.12 Electron Configurations and the
Periodic Table
Write the full electron configuration
 short hand notation
 Determine the valence electrons
 Mg ,Pd, Br
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5.15 Atomic Size
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A. Periodicity is the
presence of regularly
repeating pattern found in
nature
B. Atomic radius is
distance between the nuclei
of two atoms bonded
together
C. Atomic radius increases
down a group, decreases
across a period
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i. Larger n, larger size of
orbital
Examples

In each of the following sets of elements,
indicate which element has the smallest atomic
size
Ba, Ca, Ra
 P, Si, Al
 Rb, Cs, K
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Ionization Energies

Ionization energy (Ei) – the
amount of energy required to
remove the outermost
electron from an isolated
neutral atom in the gaseous
state
Examples

Which has higher ionization energy (Ei)?
 K or Br
 S or Te
 Ne or Sr