Chapter 2
The Structure of the
Atom and the
Periodic Table
2.1 Composition of the Atom
• Atom - the basic structural unit of an
element
• The smallest unit of an element that
retains the chemical properties of
that element
Electrons, Protons and Neutrons
• Atoms consist of three primary particles
• electrons
• protons
• neutrons
• Nucleus - small, dense, positively
charged region in the center of the atom
protons - positively charged particles
neutrons - uncharged particles
Characteristics of Atomic
Particles
• Electrons are negatively charged particles
located outside of the nucleus of an atom
• Protons and electrons have charges that are
equal in magnitude but opposite in sign
• A neutral atom that has no electrical
charge has the same number of protons
and electrons
• Electrons move very rapidly in a relatively
large volume of space while the nucleus is
small a dense
Symbolic Representation of
an Element
Charge of
particle
Mass
Number
Atomic
Number
A
Z
X
C
Symbol of
the atom
• Atomic number (Z) - the number of
protons in the atom
• Mass number (A) - sum of the number of
protons and neutrons
Atomic Calculations
number of protons + number of neutrons = mass number
number of neutrons = mass number - number of protons
In a neutral atom, the number of protons = number of
electrons
Very Important: the number of protons does not change!
Selected Properties of the Three Basic
Subatomic Particles
Name
Charge
Mass(amu)
Mass (grams)
5.4 x 10-4
9.1095 x 10-28
Electrons (e)
-1
Protons (p)
+1
1.00
1.6725 X 10-24
Neutrons (n)
0
1.00
1.6750 x 10-24
Determining the Composition of an
Atom
Calculate the number of protons, neutrons
and electrons in each of the following:
11
5
B
55
26
Fe
Isotopes
• Isotopes - atoms of the same element
having different masses
 contain same number of protons
4
 contain different numbers of neutrons
Isotopes of Hydrogen
Hydrogen
(Hydrogen - 1)
Deuterium
(Hydrogen - 2)
Tritium
(Hydrogen - 3)
Isotopic Calculations
• Isotopes of the same element have identical
chemical properties
• Some isotopes are radioactive
• Find chlorine on the periodic table
• What is the atomic number of chlorine?
• What is the mass given?
• This is not the mass number of an isotope
Atomic Mass
• Find chlorine (Cl) on the table. What is this
number, 35.34?
• The atomic mass - the weighted average of
the masses of all the isotopes that make up
chlorine
• Chlorine consists of chlorine-35 and
chlorine-37 in a 3:1 ratio
• Weighted average is an average corrected
by the relative amounts of each isotope
present in nature
Determining Atomic Mass
Calculate the atomic mass of naturally
occurring chlorine if 75.77% of chlorine
atoms are chlorine-35 and 24.23% of
chlorine atoms are chlorine-37
Step 1: convert the percentage to a decimal
fraction
0.7577 chlorine-35
0.2423 chlorine-37
Step 2: Multiply the decimal fraction by
the mass of that isotope to obtain the
isotope contribution to the atomic mass.
For chlorine-35:
0.7577 x 35.00 amu = 26.52 amu
For chlorine-37
0.2423 x 37.00 amu = 8.965 amu
Step 3: sum these partial weights to get the
weighted average atomic mass of chlorine:
26.52 amu + 8.965 amu = 35.49 amu
It’s just like calculating your
grade!
• 20% of the quiz average
• 60% of the exam average
• 20% of the final exam
Score = 0.2*quiz ave. + 0.6*exam ave. + 0.2*final
Atomic Mass Determination
• Nitrogen consists of two naturally occurring
isotopes
99.63% nitrogen-14 with a mass of 14.003 amu
0.37% nitrogen-15 with a mass of 15.000 amu
• What is the atomic mass of nitrogen?
Ions
• Ions - electrically charged particles that result
from a gain or loss of one or more electrons by
the parent atom
• Cation - positively charged
result from the loss of electrons
23Na  23Na+ + 1e-
• Anion - negatively charged
results from the gain of electrons
19F + 1 e-  19F-
2.2 Development of Atomic
Theory
• Dalton’s Atomic Theory - the first
experimentally based theory of atomic
structure of the atom.
Postulates of Dalton’s Atomic Theory
1. All matter consists of tiny particles
called atoms
2. An atom cannot be created, divided,
destroyed, or converted to any other
type of atom
3. Atoms of a particular element have
identical properties
4. Atoms of different elements have
different properties
5. Atoms of different elements
combine in simple whole-number
ratios to produce compounds (stable
aggregates of atoms)
6. Chemical change involves joining,
separating, or rearranging atoms
Postulates 1, 4, 5 and 6 are still regarded
as true.
Evidence for Subatomic Particles:
Electrons, Protons and Neutrons
• Electrons were the first subatomic
particles to be discovered using the
cathode ray tube
Indicated that the
particles were
negatively charged.
Evidence for Protons and Neutrons
• Protons were the next particle to be discovered, by
Rutherford
 Protons have the same size charge but opposite in sign
 Proton is 1837 times as heavy as electron
• Neutrons
 Postulated to exist in 1920’s but not demonstrated to
exist until 1932 by Chadwick.
 Almost the same mass as the proton (slightly heavier).
Thomson’s model of the atom
http://nobelprize.org/educational_games/physics
/quantised_world/structure-images/fig2b.gif
The “plum pudding” model.
Evidence for the Nucleus
• Earnest Rutherford’s “Gold Foil
Experiment” lead to the understanding of
the nucleus
 Most alpha particles pass through the foil
without being deflected
 Some particles were deflected, a few even
directly back to the source
Rutherford’s Gold Foil Experiment
• Most of the atom is empty space
• The majority of the mass is located in a
small, dense region
Rutherford’s model of the atom
http://www2.kutl.kyushu-u.ac.jp/seminar/
MicroWorld1_E/Part2_E/P25_E/atom.gif
2.3 Light, Atomic Structure, and
the Bohr Atom
• Rutherford’s atom – tiny, dense, positively
charged nucleus of protons surrounded by
electrons
• How do we describe the relationship of the
electrons to each other and the nucleus?
• The problem... our classical understanding
of physics didn’t work for the atom! This
will take some explaining.
Light and Atomic Structure
• Spectroscopy - absorption or emission of light
by atoms.
 Used to understand the electronic structure.
• To understand the electronic structure, we must
first understand light, Electromagnetic
Radiation
 travels in waves from a source
 speed of 3.0 x 108 m/s
Radio
• Knew from radio
that if we accelerate
an electron back
and forth in a wire
it will radiate a
radio wave
(electromagnetic
radiation).
Wavelengths
• Light is propagated (moves) as a collection
of sine waves
• Wavelength is the distance between identical
points on successive waves
• Each wavelength travels at the same velocity,
but has its own characteristic energy
Electromagnetic Spectrum
high energy
short wavelength
low energy
long wavelength
Bohr Theory
• Atoms can absorb and emit energy via
promotion of electrons to higher energy
levels and relaxation to lower levels
• Energy that is emitted upon relaxation is
observed as a single wavelength of light
• Spectral lines are a result of electron
transitions between allowed levels in the
atoms (in other words, allowed energies
are “quantized” in that only certain
quantities of energy are allowed.
The Bohr Atom
Electrons exist in fixed
energy levels
surrounding the nucleus
Promotion of
electron occurs as
it absorbs energy
Energy is released as
the electron travels
back to lower levels
Quantization of energy
Excited State
Relaxation
Electronic Transitions
• Amount of energy absorbed in jumping
from one energy level to a higher energy
level is a precise quantity
• Energy of that jump is the energy
difference between the orbits involved
• Orbit - what Bohr called the fixed energy
levels
• Ground state - the lowest possible energy
state
Modern Atomic Theory
• Bohr’s model of the atom when applied to
atoms with more than one electron failed to
explain their line spectra
• One major change from Bohr’s model is that
electrons do not move in orbits
• Atomic orbitals - regions in space with a
high probability of finding an electron
• Electrons move rapidly within the orbital
giving a high electron density
The Quantum Mechanical Atom
• Bohr’s model of the hydrogen atom
didn’t clearly explain the electron
structure of other atoms
 Electrons in very specific locations,
principal energy levels
 Wave properties of electrons conflict with
specific location
• Schröedinger developed equations that
took into account the particle nature and
the wave nature of the electrons
Schröedinger’s equations
• Equations that determine the probability of
finding an electron in specific region in space,
quantum mechanics
o Principal energy levels (n = 1,2,3…). n is also
known as the principal quantum number.
But there is more to the structure of the
arrangement of electrons found around
the nucleus of an atom, as we will see.
Energy Levels and Sublevels
PRINCIPAL ENERGY LEVELS
• n = 1, 2, 3, …
• The larger the value of n, the higher the energy
level and the farther away from the nucleus the
electrons are
• The number of sublevels in a principal energy
level is equal to n
 in n=1, there is one sublevel
 in n = 2, there are two sublevels
The angular momentum quantum number.
• An object, such as an electron, that moves
around another object (the nucleus in our case)
will have angular momentum.
• Given by l=0,1,...n-1
• l=0 is s subshell, l=1 is p subshell, etc.
• Subshells increase in energy:
s<p<d<f
• e.g., electrons in 3d subshell have more energy
than electrons in the 3p subshell
Sublevels in Each Energy Level
Principle energy
level (n)
Possible
subshells
1
1s
2
2s, 2p
3
3s, 3p, 3d
4
4s, 4p, 4d, 4f
Orbitals
Orbital - a specific region of a sublevel containing a
maximum of two electrons
• The number of orbitals in a subshell is given by the
magnetic quantum number.
• m = -l,...,0,...,+l
• Orbitals are named by their sublevel and principal
energy level
 1s, 2s, 3s, 2p, etc.
• Each type of orbital has a characteristic shape
 s is spherically symmetrical
 p has a shape much like a dumbbell
Orbital Shapes
• s is spherically
symmetrical
• Each p has a shape much like a dumbbell,
differing in the direction extending into space
Subshell
Number of
orbitals
s
1
p
3
d
5
f
7
m = -l, ...0,...+l
• An orbital can only hold two
electrons!! Once it is filled it
cannot accept more.
Electron Configuration
• Electron Configuration - the
arrangement of electrons in atomic orbitals
• Aufbau Principle - or building up
principle helps determine the electron
configuration
 Electrons fill the lowest-energy orbital that is
available first
 Remember s<p<d<f in energy
 When the orbital contains two electrons, the
electrons are said to be paired and the
orbital is full
Rules for Writing Electron
Configurations
• Obtain the total number of electrons in the atom
from the atomic number
• Electrons in atoms occupy the lowest energy
orbitals that are available – 1s first
• No more than 2 electrons in any orbital
• Maximum number of electrons in any principal
energy level is 2(n)2
• Follow the periodic table!
Writing Electron Configurations
• H
o Hydrogen has
only 1 electron
o It is in the
lowest energy
level & lowest
orbital
o Indicate
number of
electrons with a
superscript
o 1s1
• Li
o Lithium has 3
electrons
o First two have
configuration
of Helium – 1s2
o 3rd is in the
orbital of
lowest energy
in n=2
o 1s2 2s1
Classification of Elements
According to the Type of
Subshells Being Filled
Electron Configuration Examples
• Give the complete electron
configuration of each element
Be
N
Na
Cl
Ag
Shorthand Electron
Configurations
• Uses noble gas symbols to represent the
inner shell and the outer shell or valance
shell is written after
• Aluminum- full electron configuration is:
1s22s22p63s23p1
What noble gas configuration is this?
•Neon
•Configuration is written: [Ne]3s23p1
• Remember:
o How many subshells are in each
principle energy level?
o There are n subshells in the n principle
energy level.
o How many orbitals are in each
subshell?
o s has 1, p has 3, d has 5, and f has 7
o How many electrons fit in each orbital?
o2
Shorthand Electron
Configuration Examples
N
S
Ti
Sn
2.4 The Periodic Law and the
Periodic Table
• Dmitri Mendeleev and Lothar Meyer - two
scientists working independently developed
the precursor to our modern Periodic Table.
• They noticed that as you list elements in
order of atomic mass, there is a distinct
regular variation of their properties.
• Periodic Law - the physical and chemical
properties of the elements are periodic
functions of their atomic numbers.
Classification of the Elements
Parts of the Periodic Table
• Period – a horizontal row of elements in
the periodic table. They contain 2, 8, 8,
18, 18, and 32 elements,
• Group – also called families are columns
of elements in the periodic table.
• Elements in a particular group or family
share many similarities, as in a human
family.
Category Classification of
Elements
• Metals - elements that tend to lose
electrons during chemical change,
forming positive ions.
• Nonmetals - a substance whose atoms
tend to gain electrons during chemical
change, forming negative ions.
• Metalloids - have properties intermediate
between metals and nonmetals.
Classification of Elements
• Metals:
Metals
 A substance whose atoms tend to lose
electrons during chemical change
 Elements found primarily in the left 2/3 of
the periodic table
• Properties:




High thermal and electrical conductivities
High malleability and ductility
Metallic luster
Solid at room temperature
Classification of Elements
Nonmetals
• Nonmetals:
o A substance whose atoms may gain
electrons, forming negative ions
o Elements found in the right 1/3 of the
periodic table
• Properties:
o Brittle
o Powdery solids or gases
o Opposite of metal properties
2.5 Electron Arrangement and the
Periodic Table
• The electron arrangement is the primary
factor in understanding how atoms join
together to form compounds
• Electron configuration - describes the
arrangement of electrons in atoms
• Valence electrons - outermost electrons
 The electrons involved in chemical bonding
Valence Electrons
• The number of valence electrons is
the group number for the
representative elements
• The period number gives the
energy level (n) of the valence shell
for all elements
Valence Electrons and Energy
Level
• How many valence electrons does Fluorine
have?
7 valence electrons
• What is the energy level of these electrons?
Energy level is n = 2
Determining Electron Arrangement
Practice
List the total number of electrons, total
number of valence electrons, and energy
level of the valence electrons for:
• Na
• Mg
• S
• Cl
• Ar
2.6 The Octet Rule
• The noble gases are extremely stable
 Called inert as they don’t readily bond to other
elements
• The stability is due to a full complement of
valence electrons in the outermost s and p
sublevels:
 2 electrons in the 1s of Helium
 the s and p subshells full in the outermost shell of
the other noble gases (eight electrons)
Octet of Electrons
• Elements in families other than the noble
gases are more reactive
o Strive to achieve a more stable electron
configuration
o Change the number of electrons in the atom to
result in full s and p sublevels
• Stable electron configuration is called the
“noble gas” configuration
The Octet Rule
• Octet Rule - elements usually react in such a way
as to attain the electron configuration of the noble
gas closest to them in the periodic table
o Elements on the right side of the table move right to the
next noble gas
o Elements on the left side move “backwards” to the
noble gas of the previous row
• Atoms will gain, lose or share electrons in
chemical reactions to attain this more stable
energy state
Ion Formation and the Octet Rule
• Metallic elements tend to form positively
charged ions called cations
• Metals tend to lose all their valence
electrons to obtain a configuration of the
noble gas
Na
Na+ + e-
Sodium atom
11e-, 1 valence e[Ne]3s1
Sodium ion
10e[Ne]
Ion Formation and the Octet Rule
• All atoms of a group lose the same number of
electrons
• Resulting ion has the same number of electrons as
the nearest (previous) noble gas atom
Al
Al3+ + 3e-
Aluminum atom
13e-, 3 valence e[Ne]3s23p1
Aluminum ion
10e[Ne]
Using the Octet Rule
• The octet rule is very helpful in predicting
the charges of ions in the representative
elements
• Transition metals still tend to lose electrons
to become cations but predicting the charge
is not as easy
• Transition metals often form more than one
stable ion
 Iron forming Fe2+ and Fe3+ is a common example
Examples Using the Octet Rule
• Give the charge of the
most probable ion
resulting from these
elements
Ca
Sr
S
P
• Which of the
following pairs of
atoms and ions are
isoelectronic?
Cl-, Ar
Na+, Ne
Mg2+, Na+
O2-, F-
Calculating Subatomic Particles
in Ions
• How many protons, neutrons and electrons
are in the following ions?
39
19
32
16
24
12
K

2-
S
2
Mg
2.7 Trends in the Periodic Table
• Many atomic properties correlate with
electronic structure and so also with their
position in the periodic table




atomic size
ion size
ionization energy
electron affinity
Atomic Size
• The size of an element increases moving
down from top to bottom of a group
• The valence shell is higher in energy and
farther from the nucleus traveling down the
group
• The size of an element decreases from left
to right across a period
• The increase in magnitude of positive charge in
nucleus pulls the electrons closer to the nucleus
Variation in Size of Atoms
Cation Size
Cations are smaller than their parent atom
• More protons than electrons creates an increased
nuclear charge
• Extra protons pulls the remaining electrons
closer to the nucleus
• Ions with multiple positive charges are even
smaller than the corresponding monopositive
ions
Which would be smaller, Fe2+ or Fe3+?
Fe3+
• When a cation is formed isoelectronic with a
noble gas the valence shell is lost decreasing the
diameter of the ion relative to the parent atom
Anion Size
Anions are larger than their parent
atom.
• Anions have more electrons than protons
• Excess negative charge reduces the pull
of the nucleus on each individual electron
• Ions with multiple negative charges are
even larger than the corresponding
monopositive ions
Relative Size of Select Ions and
Their Parent Atoms
Ionization Energy
• Ionization energy - The energy required to
remove an electron from an isolated atom
• The magnitude of ionization energy
correlates with the strength of the attractive
force between the nucleus and the
outermost electron
• The lower the ionization energy, the easier
it is to form a cation
ionization energy + Na  Na+ + e-
Ionization Energy of Select Elements
Electron Affinity
• Electron Affinity - The energy released
when a single electron is added to an
isolated atom
• Electron affinity gives information about
the ease of anion formation
 Large electron affinity indicates an atom
becomes more stable as it forms an anion
Br + e-  Br- + energy
Periodic Trends in Electron Affinity