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Chapter 2
The Structure of the
Atom and the
Periodic Table
Denniston
Topping
Caret
5th Edition
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
2.1 Composition of the Atom
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
2.1 Composition of the Atom
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
2.1 Composition of the Atom
Symbolic Representation of
an Element
Charge of
particle
Mass
number
A
Z
Atomic
number
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
2.1 Composition of the Atom
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
2.1 Composition of the Atom
2.1 Composition of the Atom
Atomic Composition Calculations
Calculate the number of protons, neutrons,
and electrons in each of the following:
11
5
B
55
26
Fe
2.1 Composition of the Atom
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)
2.1 Composition of the Atom
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
2.1 Composition of the Atom
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
2.1 Composition of the Atom
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
2.1 Composition of the Atom
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
2.1 Composition of the Atom
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?
2.1 Composition of the Atom
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
– 23Na  23Na+ + 1e-
• Anion - negatively charged
– results from the gain of electrons
– 19F + 1e-  19F-
2.1 Composition of the Atom
Calculating Subatomic Particles
in Ions
• How many protons, neutrons, and electrons
are in the following ions?
39
19
32
16
24
12
K
S

2-
Mg
2
2.2 Development of Atomic
Theory
• Dalton’s Atomic Theory - the first
experimentally based theory of atomic
structure of the atom
2.2 Development of Atomic
Theory
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
2.2 Development of Atomic
Theory
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.
2.2 Development of Atomic
Theory
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.
2.2 Development of Atomic
Theory
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
Radioactive Emissions
The direction taken by the radioactive emissions
indicates the presence of 3 types of emissions
2.2 Development of Atomic
Theory
6
Evidence for the Nucleus
• Initial assumed protons and neutrons were
uniformly distributed throughout the atom
• Ernest 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
2.2 Development of Atomic
Theory
Rutherford’s Gold Foil Experiment
• Most of the atom is empty space
• The majority of the mass is located in a
small, dense region
2.3 Light, Atomic Structure, and
5
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?
• Use the measurement of particle energy rather
than position
2.3 Light, Atomic Structure,
and the Bohr Atom
Models of the Atom
(a) Thomson
(b) Rutherford
2.3 Light, Atomic Structure,
and the Bohr Atom
Light
• 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
2.3 Light, Atomic Structure,
and the Bohr Atom
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
2.3 Light, Atomic Structure,
and the Bohr Atom
The emission spectrum of hydrogen lead to
the modern understanding of the
electronic structure of the atom
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
2.3 Light, Atomic Structure,
and the Bohr Atom
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
2.3 Light, Atomic Structure,
and the Bohr Atom
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
2.3 Light, Atomic Structure,
and the Bohr Atom
Bohr Theory
• Allowed levels are quantized energy levels,
orbits
• Electrons are found only in these energy levels
• Highest-energy orbits are farthest from the
nucleus
• Atoms
– Absorb energy by excitation of electrons to higher
energy levels
– Release energy by relaxation of electrons to lower
energy levels
• Energy differences may be calculated from the
wavelength of light emitted
2.3 Light, Atomic Structure,
and the Bohr Atom
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
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
2.4 The Periodic Law
and the Periodic Table
Classification of the Elements
2.4 The Periodic Law
and the Periodic Table
Important Biological Elements
2.4 The Periodic Law
and the Periodic Table
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.
2.4 The Periodic Law
and the Periodic Table
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
2.4 The Periodic Law
and the Periodic Table
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
2.4 The Periodic Law
and the Periodic Table
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
2.4 The Periodic Law
and the Periodic Table
Classification of Elements
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
2.4 The Periodic Law
and the Periodic Table
Classification of Elements
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
2.4 The Periodic Law
and the Periodic Table
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
2.4 The Periodic Law
and the Periodic Table
Element Information in the
Periodic Table
20
Ca
Calcium
40.08
atomic number
symbol
name
atomic mass
2.4 The Periodic Law
and the Periodic Table
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?
2.5 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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
Electron Arrangement by
Energy Level
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2. Atomic number = number of electrons in
an atom
•
Silicon has 14 electrons
2.5 Electron Arrangement
and the Periodic Table
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
4. Energy levels:
•
•
•
n = 1 holds 2 electrons
n = 2 holds 8 electrons (total of 10)
n = 3 holds remaining 4 electrons (total = 14)
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
Schröedinger’s equations
• Equations that determine the probability
of finding an electron in specific region
in space, quantum mechanics
– Principle energy levels (n = 1, 2, 3…)
– Each energy level has one or more
sublevels or subshells (s, p, d, f)
– Each sublevel contains one or more
atomic orbitals
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
Orbital Shapes
• s is spherically
symmetrical
• Each p has a shape much like a dumbbell,
differing in the direction extending into space
2.5 Electron Arrangement
and the Periodic Table
Subshell
Number of
orbitals
s
1
p
3
d
5
f
7
• How many electrons can be in the
4d subshell?
•10
• 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
Shell 4
4f •• •• •• •• •• •• ••
Increasing Energy
2.5 Electron Arrangement
and the Periodic Table
Quantum Mechanical Model
4d
•• •• •• •• ••
Sublevel
4p
•• •• ••
Orbital
4s
••
Electron
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
Electron Filling Order
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
Orbital Energy-level Diagram
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
Electron Configuration Examples
• Give the complete electron
configuration of each element
– Be
–N
– Na
– Cl
– Ag
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
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
2.5 Electron Arrangement
and the Periodic Table
• 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
2.5 Electron Arrangement
and the Periodic Table
Shorthand Electron
Configuration Examples
• N
• S
• Ti
• Sn
2.5 Electron Arrangement
and the Periodic Table
Classification of Elements
According to the Type of
Subshells Being Filled
Use this breakdown of the Periodic Table and you can
write the configuration of any element.
2.5 Electron Arrangement
and the Periodic Table
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 highestenergy 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
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)
2.6 The Octet Rule
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
2.6 The Octet Rule
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
2.6 The Octet Rule
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]
2.6 The Octet Rule
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]
2.6 The Octet Rule
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
O + 2e-
O2-
Oxygen atom
8e-, 6 valence e[He]2s22p4
Oxide ion
10e[He]2s22p6 or [Ne]
2.6 The Octet Rule
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
2.6 The Octet Rule
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-
2.7 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
2.7 Trends in the Periodic
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
2.7 Trends in the Periodic
Table
Variation in Size of Atoms
2.7 Trends in the Periodic
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
2.7 Trends in the Periodic
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
2.7 Trends in the Periodic
Table
Relative Size of Select Ions and
Their Parent Atoms
2.7 Trends in the Periodic
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
ionization energy + Na  Na+ + e-
2.7 Trends in the Periodic
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?
2.7 Trends in the Periodic
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
Br + e–  Br– + energy
2.7 Trends in the Periodic
Table
Periodic Trends in Electron
Affinity
• Electron affinity
generally
decreases down a
group
• Electron affinity
generally increases
across a period