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Unit 2
Quantum Theory, Electrons,
& The Periodic Table
Chapters 4 & 5
Arrangement of Electrons in Atoms
CHAPTER 4
Chapter 4 – Section 1: The Development of a New Atomic Model
Properties of Light
• Sometimes light behaves like waves, and
other times like particles.
• Visible light is a kind of electromagnetic
radiation, which is a form of energy that
exhibits wavelike
behavior as it
travels through
space.
Chapter 4 – Section 1: The Development of a New Atomic Model
The Electromagnetic Spectrum
• Together, all the forms of electromagnetic
radiation form the electromagnetic spectrum.
Visual Concept
Chapter 4 – Section 1: The Development of a New Atomic Model
Visible Light
• Visible Light is the narrow band of
electromagnetic radiation that we can see.
Visible Spectrum
• It consists of a
range of waves
Color
Wavelength
Red
700 - 650 nm
with various
Orange
649 - 580 nm
Yellow
579 - 575 nm
wavelengths.
Green
Blue
Indigo
Violet
574 - 490 nm
489 - 455 nm
454 - 425 nm
424 - 400 nm
Chapter 4 – Section 1: The Development of a New Atomic Model
The Speed of Light
• The constant, c, equals the speed of light,
and it is a fundamental constant of the
universe.
• All waves in the
electromagnetic
spectrum travel at
the speed of light,
c = 3 x 108 m/s.
Chapter 4 – Section 1: The Development of a New Atomic Model
Properties of Light (continued)
• Wavelength (λ) is the distance between
corresponding points
on adjacent waves.
• Frequency (ν)
is defined as the
number of waves
that pass a given
point in a specific
time, usually
one second.
Chapter 4 – Section 1: The Development of a New Atomic Model
Wavelength vs. Frequency
• Wavelength (λ) is inversely proportional to
frequency (ν). In other words, when λ
increases, ν decreases,and vice versa.
Chapter 4 – Section 1: The Development of a New Atomic Model
Wavelength vs. Frequency (continued)
• The relationship between wavelength and
frequency is described by the equation:
c = λν
• Where c is a constant (always the same
number) equal to 3x108 m/s.
• λ is the wavelength (in m).
Problem-solving hint: 1 nm = 10-9 m.
• ν is the frequency (in s−1 or Hz).
Visual Concept
Chapter 4 – Section 1: The Development of a New Atomic Model
The Speed of Light
Sample Problem
A photon of light has a frequency of 4.4x1014 Hz.
Calculate its wavelength. Does it fall within the
visible spectrum? If so, what color is it?
Solution:
Visible Spectrum
Color
Wavelength
Use the equation: c = λν
Red
700 - 650 nm
Orange
649 - 580 nm
3x108m/s = λ (4.4x1014Hz)
Yellow
579 - 575 nm
Green
574 - 490 nm
3 x 108m/s
489 - 455 nm
-7m Blue
λ=
=
6.8
x10
Indigo
454 - 425 nm
14
4.4 x 10 Hz
Violet
424 - 400 nm
λ = 680 x 10-9 m, or
λ = 680 nm
Yes, it is red light
Chapter 4 – Section 1: The Development of a New Atomic Model
The Photoelectric Effect
• By the early 1900s, scientists observed
interactions of light and matter that couldn’t
be explained by wave theory.
• The photoelectric effect
refers to the emission of
electrons from a metal
when light shines on
the metal.
Visual Concept
Chapter 4 – Section 1: The Development of a New Atomic Model
The Particle Description of Light
• A quantum of energy is the minimum quantity
of energy that can be lost or gained by an atom.
• A photon is a particle of
electromagnetic radiation
having zero mass and
carrying a quantum of energy.
• The energy of a particular photon is directly
proportional to the frequency of the radiation.
Visual Concept
Chapter 4 – Section 1: The Development of a New Atomic Model
Energy States
• Ground state – The lowest
energy state of an atom.
• Excited state – an atom has
a higher potential energy
than it has in its ground state.
• When an excited atom returns
to its ground state, it gives off
energy in the form of
electromagnetic radiation.
Chapter 4 – Section 1: The Development of a New Atomic Model
Hydrogen’s Line Emission Spectrum
Chapter 4 – Section 1: The Development of a New Atomic Model
The Bohr Model
• In 1913, Danish physicist Niels Bohr
proposed a hydrogen-atom model
that linked the atom’s electron to
its line-emission spectrum.
Chapter 4 – Section 1: The Development of a New Atomic Model
The Bohr Model (continued)
• According to the Bohr model, the electron
can circle the nucleus only in allowed
paths, or orbits.
• The energy of the
electron is higher
when it is in orbits
that are farther
from the nucleus.
Visual Concept
Chapter 4 – Section 2: The Quantum Model of the Atom
Electrons as Waves
• In 1924, French scientist Louis de
Broglie suggested that electrons
act like waves confined to the
space around an atomic nucleus.
• It followed that the electron waves could
exist only at specific
frequencies –
corresponding to the
quantized energies
of Bohr’s orbits.
Chapter 4 – Section 2: The Quantum Model of the Atom
The Heisenberg Uncertainty Principle
• In 1927, German physicist Werner
Heisenberg realized that an attempt
to locate an electron with a photon
knocks the electron off its course.
• The Heisenberg uncertainty
principle states that it is impossible
to determine simultaneously both the
position and velocity of an electron
or any other very small particle.
Visual Concept
Chapter 4 – Section 2: The Quantum Model of the Atom
The Schrödinger Wave Equation
• In 1926, Austrian physicist
Erwin Schrödinger developed
an equation that treated
electrons in atoms as waves.
• Together with Heisenberg
and others, Schrödinger laid
the foundation for modern
quantum theory.
Visual Concept
Chapter 4 – Section 2: The Quantum Model of the Atom
Quantum Theory
• Quantum theory describes
mathematically the wave
properties of electrons and
other very small particles.
• There are four different
types of quantum numbers used:
1.
2.
3.
4.
Principal quantum # (n) – energy level.
Angular momentum quantum # (l) - sublevel.
Magnetic quantum # (m) - orbital.
Spin quantum # (s).
Chapter 4 – Section 2: The Quantum Model of the Atom
Principal Quantum Number
• The Principal Quantum Number (n)
indicates the main energy level occupied by
an electron.
• As n increases, the electron’s energy and
its distance from the
nucleus increases.
Chapter 4 – Section 2: The Quantum Model of the Atom
Angular Momentum Quantum Number
• The Angular Momentum Quantum
Number (l) (also called the sublevel)
indicates the shape of the orbital.
• The number of sublevels allowed
for each energy
level is equal
s orbital
to n.
p orbital
sphere
dumbbell
d orbital
cloverleaf
f orbital
complex
Chapter 4 – Section 2: The Quantum Model of the Atom
Magnetic Quantum Number
• The Magnetic Quantum Number (m)
indicates the orientation of an orbital around
the nucleus.
Sublevel Orbitals
s
1
p
3
d
5
f
7
Chapter 4 – Section 2: The Quantum Model of the Atom
Spin Quantum Number
• The Spin Quantum Number (s) indicates
the fundamental spin state of an electron in
an orbital.
• There are only two possible
values for s, +½ and –½.
• A single orbital can hold a maximum of two
electrons, but the
electrons must have
opposite spin states.
Chapter 4 – Section 2: The Quantum Model of the Atom
Quantum Numbers Overview
Chapter 4 – Section 3: Electron Configuration
Electron Configuration Rules
• According to the Aufbau Principle, an
electron occupies the lowest-energy orbital
that can receive it.
• The order of
increasing energy
is shown on the
vertical axis. Each
box represents an
orbital. (diagram
on pg. 111)
Chapter 4 – Section 3: Electron Configuration
Electron Configuration Rules (continued)
• According to the
Pauli exclusion
principle, no two
electrons in the
same atom can
have the same set
of four quantum numbers.
Chapter 4 – Section 3: Electron Configuration
Electron Configuration Rules (continued)
• According to Hund’s Rule,
orbitals of equal energy are
each occupied by one
electron before any orbital
is occupied by a second
electron, and all electrons
in singly occupied orbitals
must have the
same spin state.
Wrong
Wrong
Correct
Chapter 4 – Section 3: Electron Configuration
Orbital Notation
• An orbital containing one electron is
represented as:
• An orbital containing two electrons is
represented as:
• The lines are labeled with the principal
quantum number and sublevel letter. For
example, the orbital notation for helium is
written as follows:
He
1s
Chapter 4 – Section 3: Electron Configuration
Orbital Notation
Sample Problem 1
Write the orbital notation for Carbon.
Solution:
Carbon is atomic number 6, so it has 6 electrons.
The first two electrons go in the 1s orbital.
The next two electrons go in the 2s orbital.
The final two electrons go in the 2p orbitals.
Carbon
1s
2s
2p
Chapter 4 – Section 3: Electron Configuration
Electron Configuration Notation
• Electron-configuration notation eliminates
the lines and arrows of orbital notation.
• Instead, the number of electrons in a
sublevel is shown by a superscript.
• Example: Carbon
Orbital Notation
Electron Configuration
1s22s22p2
1s
2s
2p
Visual Concept
Chapter 4 – Section 3: Electron Configuration
Blocks of the Periodic Table
s
p
d
f
Chapter 4 – Section 3: Electron Configuration
Electron Configuration Notation
Sample Problem 1
a. Write electron configuration for Selenium (Se).
b. How many unpaired electrons are in an atom
of Selenium?
Solution:
a. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
b. Only consider the 4p4 electrons, since all
electrons will be paired in filled orbitals.
2 electrons are unpaired
4p
Chapter 4 – Section 3: Electron Configuration
Noble Gas Notation
• The Group 18 elements (He, Ne, Ar, Kr,
Xe, and Rn) are called the noble gases.
• Noble gas notation is an abbreviated
electron configuration.
• Use square brackets around the noble gas
at the end of the prior period to replace
part of the configuration.
• Example: Calcium
Electron Configuration
1s22s22p63s23p64s2
Noble Gas Notation
[Ar]4s2
Chapter 4 – Section 3: Electron Configuration
Noble Gas Notation
Sample Problem 1
a. Write the noble gas notation for Gold (Au).
b. How many inner-shell electrons does this
atom have?
Solution:
a. [Xe] 6s2 4f14 5d9
b. The outer shell is the one with the highest #.
There are 2 e- in energy level 6 (6s2).
All the rest are inner-shell electrons.
79 total e- - 2 outer-shell e- = 77 inner-shell e-
The Periodic Law
CHAPTER 5
Chapter 5 – Section 1: History of the Periodic Table
Mendeleev and Periodicity
• The first periodic table of the
elements was published in 1869 by
Russian chemist Dmitri Mendeleev.
• Mendeleev left empty spaces in his
table and predicted
elements that would fill
3 of the spaces.
• By 1886, all 3 of these
elements had been
discovered.
Visual Concept
Chapter 5 – Section 1: History of the Periodic Table
Mosley and the Periodic Law
• In 1911, the English scientist
Henry Moseley discovered that
the elements fit into patterns
better when they were arranged
according to atomic number,
rather than atomic weight.
• The Periodic Law states that the physical
and chemical properties of the elements are
periodic functions of their atomic numbers.
Chapter 5 – Section 2: Electron Configuration and the Periodic Table
The Periodic Table
• Elements in the periodic table are arranged
into vertical columns, called groups or
families, that share similar chemical
properties.
• Elements are
also organized
horizontally
in rows,
or periods.
Chapter 5 – Section 2: Electron Configuration and the Periodic Table
Group 1: Alkali Metals
• Group 1 elements are called alkali metals.
• Alkali metals have a silvery appearance
and are soft enough to cut with a knife.
• They are extremely reactive and are not
found in nature as free elements.
• They must be stored under oil or kerosene.
Visual Concept
Chapter 5 – Section 2: Electron Configuration and the Periodic Table
Group 2: Alkaline Earth Metals
• Elements in group 2 are known
as the alkaline earth metals.
• Group 2 metals are harder, denser
and stronger than alkali metals,
and have higher melting points.
• Less reactive than
group 1, but still too
reactive to be found
in nature as free
elements.
Chapter 5 – Section 2: Electron Configuration and the Periodic Table
Group 17: Halogens
• Elements in group 17 are
known as the halogens.
• Halogens are the most
reactive nonmetals,
reacting vigorously with
metals to form salts
•Most halogens exist in
nature as diatomic molecules
(i.e. F2, Cl2, Br2 and I2.)
Visual Concept
Chapter 5 – Section 2: Electron Configuration and the Periodic Table
Group 18: Noble Gases
• Elements in group 18 are
known as noble gases.
• They are completely nonreactive and don’t form
compounds under normal
conditions.
• A new group was added to
the periodic table in 1898
for the noble gases.
Chapter 5 – Section 2: Electron Configuration and the Periodic Table
d-block: Transition Metals
• Elements in the d-block are
called transition metals.
• They have typical metallic
properties such as conduction
of electricity and high luster.
• Less reactive than group 1 and 2 elements.
• Some (i.e. platinum & gold) are so unreactive
they usually don’t form compounds.
Chapter 5 – Section 2: Electron Configuration and the Periodic Table
f-block: Lanthanides & Actinides
• Elements in the period 6
of the f-block are called
lanthanides (or rare-earth).
• Lanthanides are shiny
metals similar in reactivity
to alkaline earth metals.
• Elements in period 7 of the
f-block are called actinides.
• Actinides are all radioactive,
and many of them are known
only as man-made elements.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Atomic Radii
• Atomic radius – one-half the distance
between the nuclei of identical atoms
Group 1
that are bonded together.
•Atomic radii tend to increase
as you go down a group
because electrons occupy
successively higher energy
levels farther away from the
nucleus.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Atomic Radii (continued)
• Atomic radii tend to decrease as you go
across a period because as more
electrons are added they are pulled closer
to the more highly charged nucleus.
Period 2
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Atomic Radii
Sample Problem
Of the elements Mg, Cl, Na, and P, which has
the largest atomic radius? Explain.
Solution:
Na has the
largest radius.
All of the
elements are
in the 3rd period,
and atomic radii
decrease across
a period.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Ionization Energy
• An ion is an atom of group of bonded atoms
that has a positive or negative charge.
• The energy required to
remove an electron from
a neutral atom of an
element is called the ionization energy (IE).
• Ionization energy tends to increase across
each period because a higher nuclear
charge more strongly attracts electrons in
the same energy level.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Ionization Energy (continued)
• Ionization energy tends to decrease down
each group because electrons farther from
the nucleus are removed more easily.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Ionization Energy
Sample Problem
Consider two elements, A and B. A has an IE of 419
kJ/mol. B has an IE of 1000 kJ/mol. Which element is
more likely to be in the s block? Which will be in the
p block? Which is more likely to form a positive ion?
Solution:
Element A is most likely to be in the s-block since IE
increases across the periods.
Element B would most likely lie at the end of a period
in the p block.
Element A is more likely to form a positive ion since it
has a much lower IE than B.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Electron Affinity and Electronegativity
• Electron affinity is the energy change that
occurs when an electron is acquired by a
neutral atom.
• Electronegativity is a measure of the ability
of an atom in a chemical compound to
attract electrons from another atom in the
compound.
• Electronegativity applies to atoms in a
compound, while electron affinity is a
property of isolated atoms.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Electron Affinity and Electronegativity
(continued)
• Electron affinity and electronegativity both
tend to increase across periods, and
decrease (or stay the same) down a group.
Chapter 5 – Section 3: Electron Configuration and Periodic Properties
Electronegativity
Sample Problem
Of the elements Ga, Br, and Ca, which has
the highest electronegativity? Explain .
Solution:
All of these elements are in the fourth period.
Br has the highest atomic number and is
farthest to the right in the period.
Br would have the highest electronegativity
since electronegativity increases across
a period.