Structure of Atoms and Periodic Table

Chapter 2
The Structure of the
Atom and the Periodic
Table
Denniston
Topping
Caret
7th
Edition
Copyright© The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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

2.1CompositionoftheAtom
• Nucleus - small, dense, positively
charged region in the center of the atom
- protons - positively charged particles
- neutrons - uncharged particles
Electrons, Protons, and Neutrons
• Atoms consist of three primary particles
• electrons
• protons
• neutrons

2.1CompositionoftheAtom 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 and dense

Mass
Atomic
number
Charge of
particle
Symbol of
the atom
2.1CompositionoftheAtom Symbolic Representation of
an Element
CA
Z X
• 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
number of protons = number of electrons IF positive and
negative charges cancel, the atom charge = 0

Calculate the number of protons, neutrons,
and electrons in each of the following:
B11
5
Fe55
26
Atomic Composition Calculations

4
Hydrogen
(Hydrogen - 1)
Deuterium
(Hydrogen - 2)
Tritium
(Hydrogen - 3)
Isotopes of Hydrogen
• Isotopes - atoms of the same element
having different masses
– contain same number of protons
– contain different numbers of neutrons
Isotopes

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?
17
• What is the mass given?
35.45
• This is not the mass number of an isotope

2.1CompositionoftheAtom Atomic Mass
• 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

Atomic Mass Calculation
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

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 and Charges
• Ions - electrically charged particles that
result from a gain or loss of one or more
electrons by the parent atom
• Cation - positively charged
– results from the loss of electrons
– 23
Na  23
Na+
+ 1e-
• Anion - negatively charged
– results from the gain of electrons
– 19
F + 1e-
 19
F-

+
K39
19
-232
16S
+224
12 Mg
2.1CompositionoftheAtom Calculating Subatomic Particles
in Ions
• How many protons, neutrons, and electrons
are in the following ions?

2.2 Development of Atomic
Theory
• Dalton’s Atomic Theory - the first
experimentally based theory of atomic
structure of the atom

2.2Developmentof
AtomicTheoryPostulates 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.
2.2Developmentof
AtomicTheory

• Electrons were the first subatomic
particles to be discovered using the
cathode ray tube.
Indicated that the
particles were
negatively charged.
2.2Developmentof
AtomicTheory
Subatomic Particles:
Electrons, Protons, and Neutrons

2.2Developmentof
AtomicTheory Evidence for Protons and
Neutrons
• Protons were the next particle to be discovered,
by Goldstein
– Protons have the same size charge but opposite in sign
– A proton is 1,837 times as heavy as an electron
• Neutrons
– Postulated to exist in 1920’s but not demonstrated to
exist until 1932
– Almost the same mass as the proton

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 Elements2.4ThePeriodicLaw
andthePeriodicTable

Important Biological Elements2.4ThePeriodicLaw
andthePeriodicTable

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, and are
columns of elements in the periodic table.
• Elements in a particular group or family
share many similarities, as in a human
family.
2.4ThePeriodicLaw
andthePeriodicTable

2.4ThePeriodicLaw
andthePeriodicTable
Families of the Periodic Table
• Representative elements - Group A
elements
• Transition elements - Group B
elements
• Alkali metals - Group IA
• Alkaline earth metals - group IIA
• Halogens - group VIIA
• Noble gases - group VIIIA

2.4ThePeriodicLaw
andthePeriodicTable
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
2.4ThePeriodicLaw
andthePeriodicTable

Nonmetals
• Nonmetals:
– A substance whose atoms may gain
electrons, forming negative ions
– Elements found in the right 1/3 of the
periodic table
• Properties:
– Brittle
– Powdery solids or gases
– Opposite of metal properties
2.4ThePeriodicLaw
andthePeriodicTable

Metalloids
• Metalloids:
– Elements that form a narrow diagonal band
in the periodic table between metals and
nonmetals
• Properties are somewhat between those
of metals and nonmetals
• Also called semimetals
2.4ThePeriodicLaw
andthePeriodicTable

2.4ThePeriodicLaw
andthePeriodicTable
Atomic Number and Atomic Mass
• Atomic Number:
– The number of protons in the nucleus of
an atom of an element
– Nuclear charge or positive charge from
the nucleus
• Most periodic tables give the element
symbol, atomic number, and atomic
mass

2.4ThePeriodicLaw
andthePeriodicTable
Element Information in the
Periodic Table
20 atomic number
Ca symbol
Calcium name
40.08 atomic mass

Using the Periodic Table
• Identify the group and period to
which each of the following belongs:
a. P
b. Cr
c. Element 30
• How many elements are found in
period 6?
• How many elements are in group
VA?
2.4ThePeriodicLaw
andthePeriodicTable

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

2.5ElectronArrangement
andthePeriodicTable 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
andthePeriodicTable

Electron Arrangement by
Energy Level
andthePeriodicTable

andthePeriodicTable
Valence Electrons - Detail
• What is the total number of electrons in
fluorine?
– Atomic number = 9
– 9 protons and 9 electrons
• 7 electrons in the valence shell, (n = 2 energy level),
so where are the other two electrons?
– In n = 1 energy level
– Level n = 1 holds only two electrons

Determining Electron Arrangement
List the total number of electrons, total number of
valence electrons, and energy level of the valence
electrons for silicon.
1. Find silicon in the periodic table
• Group IVA
• Period 3
• Atomic number = 14
1. Atomic number = number of electrons
in an atom
• Silicon has 14 electrons
andthePeriodicTable

Determining Electron Arrangement #2
List the total number of electrons, total number of
valence electrons, and energy level of the valence
electrons for silicon.
3. As silicon is in Group IV, only 4 of its 14
electrons are valence electrons
• Group IVA = number of valence electrons
3. Energy levels:
• n = 1 holds 2 electrons
• n = 2 holds 8 electrons (total of 10)
• n = 3 holds remaining 4 electrons (total = 14)
andthePeriodicTable

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
andthePeriodicTable

andthePeriodicTable
Energy Levels and Subshells
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

andthePeriodicTable
Principal Energy Levels
• The electron capacity of a principal
energy level (or total electrons it can hold) is
2(n)2
– n = 1 can hold 2(1)2
= 2 electrons
– n = 2 can hold 2(2)2
= 8 electrons
• How many electrons can be in the n = 3
level?
– 2(3)2
= 18
• Compare the formula with periodic table…..

n = 1, 2(1)2
= 2
n = 2, 2(2)2
= 8
n = 3, 2(3)2
= 18
n = 4, 2(4)2
= 32

andthePeriodicTable
Sublevels
• Sublevel: a set of energy-equal orbitals
within a principal energy level
• Subshells increase in energy:
s < p < d < f
• Electrons in 3d subshell have more energy
than electrons in the 3p subshell
• Specify both the principal energy level and a
subshell when describing the location of an
electron

Principle energy
level (n)
Possible
subshells
1 1s
2 2s, 2p
3 3s, 3p, 3d
4 4s, 4p, 4d, 4f
andthePeriodicTable
Sublevels in Each Energy Level

andthePeriodicTable
Orbitals
• Orbital - a specific region of a sublevel
containing a maximum of two electrons
• 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

andthePeriodicTable
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
• How many electrons can be in the
4d subshell?
•10
andthePeriodicTable

Quantum Mechanical Model
• Each orbital within a
sublevel contains a
maximum of 2
electrons
• Energy increases as n,
shell number
increases, but ALSO
increases as you move
from s to p to d to f
sublevels
andthePeriodicTable
IncreasingEnergy 4s
4p
4d
4f
••
•• •• ••
•• •• •• •• ••
••••••••••••••
Electron
Orbital
Sublevel
Shell 4

andthePeriodicTable
Electron Spin
• 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

andthePeriodicTable
Electron Filling Order

andthePeriodicTable
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
• Each principal energy level, n contains only n
sublevels
• Each sublevel is composed of orbitals
• No more than 2 electrons in any orbital
• Maximum number of electrons in any principal
energy level is 2(n)2

Electron Distribution
• This table lists the number of electrons in each
shell for the first 20 elements
• Note that 3rd
shell stops filling at 8 electrons even though
it could hold more
andthePeriodicTable

Orbital Energy-level Diagram
andthePeriodicTable

andthePeriodicTable
Writing Electron Configurations
• H
– Hydrogen has
only 1 electron
– It is in the
lowest energy
level & lowest
orbital
– Indicate
number of
electrons with a
superscript
– 1s1
• Li
– Lithium has 3
electrons
– First two have
configuration
of Helium – 1s2
– 3rd
is in the
orbital of
lowest energy
in n=2
– 1s2
2s1

andthePeriodicTable
Electron Configuration Examples
• Give the complete electron
configuration of each element
– Be
– N
– Na
– Cl
– Ag

The Shell Model and Chemical
Properties
• As we explore the model placing electrons
in shells, we will see that the pattern which
emerges from this placement correlates well
with a pattern for various chemical
properties
• We will see that all elements in a group
have the same number of electrons in their
outermost (or valence) shell
andthePeriodicTable

Groups Have Similar Chemical
Properties and Appearances
• Examples of different elements that
have similar properties and are all in
group VA
– Nitrogen
– Phosphorus
– Arsenic
– Antimony
– Bismuth
andthePeriodicTable

What noble gas configuration is this?
•Neon
•Configuration is written: [Ne]3s2
3p1
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:
1s2
2s2
2p6
3s2
3p1
andthePeriodicTable

• Remember:
– How many subshells are in each
principle energy level?
– There are n subshells in the n principle
energy level.
– How many orbitals are in each
subshell?
– s has 1, p has 3, d has 5, and f has 7
– How many electrons fit in each orbital?
– 2
andthePeriodicTable

andthePeriodicTable
Shorthand Electron
Configuration Examples
• N
• S
• Ti
• Sn

Use this breakdown of the Periodic Table and you can
write the configuration of any element.
andthePeriodicTable
According to the Type of
Subshells Being Filled

Classification of Elements –
by Group
• Representative element: An element in which the
distinguishing electron is found in an s or p
subshell
• Distinguishing electron: The last or highest-
energy electron found in an element
• Transition element: An element in which the
distinguishing electron is found in a d subshell
• Inner-transition element: An element in which
the distinguishing electron is found in a f
subshell
andthePeriodicTable

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 are 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
– Strive to achieve a more stable electron
configuration
– 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
2.6TheOctetRule

2.6TheOctetRule 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
– Elements on the right side of the table move right to the
next noble gas
– 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

2.6TheOctetRule
Na
Sodium atom
11e-
, 1 valence e-
[Ne]3s1
Na+
+ e-
Sodium ion
10e-
[Ne]
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

2.6TheOctetRule
Al
Aluminum atom
13e-
, 3 valence e-
[Ne]3s2
3p1
Al3+
+ 3e-
Aluminum ion
10e-
[Ne]
• 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
Ion Formation and the Octet Rule

O + 2e-
Oxygen atom
8e-
, 6 valence e-
[He]2s2
2p4
O2-
Oxide ion
10e-
[He]2s2
2p6
or [Ne]
2.6TheOctetRule Isoelectronic
• Isoelectronic - atoms of different elements having
the same electron configuration (same number of
electrons)
• Nonmetallic elements, located on the right side of
the periodic table, tend to form negatively charged
ions called anions
• Nonmetals tend to gain electrons so they become
isoelectronic with its nearest noble gas neighbor
located in the same period to the right

2.6TheOctetRule 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-
2.6TheOctetRule

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

2.7TrendsinthePeriodic
Table 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

Table Variation in Size of Atoms

Table
Cation Size
Cations are smaller than their parent atom
• More protons than electrons creates an increased
nuclear charge
• Extra protons pull 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

Table
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

TableRelative Size of Select Ions and
Their Parent Atoms

ionization energy + Na  Na+
+ e-
Table
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

Table
Ionization Energy of Select Elements
• Ionization decreases down a family as the
outermost electrons are farther from the nucleus
• Ionization increases across a period because the
outermost electrons are more tightly held
• Why would the noble gases be so unreactive?

Br + e–
 Br–
+ energy
Table
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

Table
Periodic Trends in Electron
Affinity
• Electron affinity
generally
decreases down a
group
• Electron affinity
generally increases
across a period

Structure of Atoms and Periodic Table

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