Chapter 4
Atoms and
Elements
Homework
 Assigned
Problems (odd numbers only)
 “Questions
and Problems” 4.1 to 4.65
(begins on page 96)
 “Additional
Questions and Problems”
4.75 to 4.105 (page 122-123)
 “Challenge
(page 124)
Questions” 4.107 to 4.111,
Elements
Matter is anything with a mass and occupies space
 Matter (in our world) is composed of combinations
of about 100 basic substances called elements
 109 elements have been discovered and isolated
 88 are found in nature
 21 are (synthetic) man-made
 Oxygen most abundant element (by mass) on earth

“Element”

A pure substance that cannot be broken down into
simpler substances by a chemical means
 Single atom of that element
 Sample of the element large enough to weigh on a
scale
 Generally referring to the presence of that element
(compound), not necessarily in its free form
Chemical symbols
 Each element has a unique symbol
 One or two letter abbreviations
 If
two letters, the second is lower case
 The
letter symbol often corresponds to
the name of the element
 F = Fluorine
 P = Phosphorous
 Some
symbols derived from the Latin or
Greek names
 Lead – Pb (plumbum)
 Gold – Au (aurum)
 Sodium – Na (natrium)
Names of Symbols and Some Common
Elements
 Required: Know the name and
symbol of some of the most
common elements
 Table 4.2 on p. 95 (know the
names and symbols listed)
 A periodic table will be given on
each test or quiz
 Required: Know how to use a
periodic table to find needed
information
Periodic Table
 A chart
of the elements with similar
chemical properties arranged into vertical
columns (groups)
 Horizontal rows are called periods
 The elements arranged (in rows) in order
of increasing atomic mass (atomic
number)
 Main group elements are those in the
columns labeled with numbers (1A-8A)
 Transition elements are those in the
columns labeled with the letter “B”
Main Group Elements
 Also
called representative elements
 The elements in the A-groups
 First two columns (1A and 2A)
 The last 6 columns (3A to 8A)
 Easy to predict ionic structure
Transition Elements
The
elements in the B-groups
Middle block of elements (3B
through 2B)
Includes the two groups at the
bottom
Lanthanides
Difficult
and Actinides
to predict ionic structure
Classification of Elements

Certain groups of elements have their own
special names due to the chemical similarity of
the elements in them
Alkali Metals
 Group 1A
 Alkaline Earth Metals
 Group 2A
 Halogens
 Group 7A
 Noble Gases
 Group 8A
 Become familiar with these group names

The Periodic Table
1A
1
2
3
4
5
6
7
2A
8A
3A 4A 5A 6A7A
Metals/Nonmetals
 Metals
Everything
to the left of the
metal/nonmetal barrier
Shiny solid, good conductor of
electricity, ductile and malleable
 Nonmetals
Everything to the right of the
metal/nonmetal barrier
Dull appearance, not ductile or
malleable not good conductors of
electricity
Metalloids
Elements with properties
intermediate between those
metals and nonmetals
On the metal/nonmetal barrier
Have some physical properties
of metals but some chemical
properties of nonmetals
Semiconductors
Si,
Ge, As, Sb, Te
The Atom
The smallest particle of an
element that can exist and still
have properties of that element
All atoms of a certain type are
similar to one another and
different from all other types
109 different types are known
and each “type” is a different
element
Dalton’s Atomic Theory (1808)
 A set of five statements that
summarize the modern
scientific concept about atoms
1) All matter is made from small
particles called atoms (109
different types)
2) All atoms of a given type are
similar to one another and
significantly different from all
other types
`
Dalton’s Atomic Theory
3) The number and arrangement of
different types of atoms in a pure
substance determine its identity
4) A chemical change is a
combination, separation, or
rearrangement of atoms (forms
new substances)
5) Only whole atoms take part or
result from any chemical reaction
`
Dalton’s Atomic Theory
 Atoms
are indivisible in a chemical
process (indestructible)
All atoms present at beginning are
present at the end
Atoms are not created or
destroyed, just rearranged
Atoms of one element cannot
change into atoms of another
element
Cannot
turn Lead into Gold by a
chemical reaction
Cathode Rays and Electrons

J.J. Thomson (1897) used a gas discharge tube to
investigate a beam called a cathode ray
 Determined that the ray was made of tiny negatively
charged particles we call electrons
 He determined the electrons were smaller than a
hydrogen atom
 Since electrons are smaller than atoms they must be
parts of an atom
 Atoms must be divisible
 Atoms of different elements all produced these same
electrons
Parts of an Atom: The Electron
Defined
by Thompson
Tiny, negatively charged particle
Charge
Very
is -1
light compared to mass of
atom
1/2000th
Moves
atom
the mass of a H atom
very rapidly within the
Thompson’s Model of Atoms
 Atoms have a structure and are divisible
 Thomson reasoned that electrons must
be a fraction of the entire size of the
atom since their mass is much smaller
that the whole atom
 Thomson also reasoned since atoms are
neutral, the electrons were embedded in
a sphere of uniform positive charge
 Thomson (1898) proposed the “Plum
Pudding” model or “Raisin Muffin” model
of the atom
Thompson’s Model of Atoms

Thomson Atomic
Model (early 1900’s):
Proposed a uniform,
positive sphere of
matter with small
negative electrons
attached to the
surface of the sphere
 This became known
as the plum-pudding
model
Rutherford’s Experiment
 1911
Rutherford designed an experiment
to test the Thompson model (“plumpudding”) of the atom
 Rutherford
directed positively charged
particles (alpha particles) towards a thin
gold foil sheet
 Rutherford
expected the particles to pass
straight through a uniform area of mass
and positive charge
Rutherford’s Experiment
Some 
particles are
scattered
Most particles
pass straight
through foil
Source of
 particles
Beam of
 particles
Fig4_5
Screen to detect
scattered  particles
Thin
metal foil
Rutherford’s Experiment
Results:
 Most (alpha) particles
mostly went straight
through
 A few particles were
unexpectedly deflected
from their expected
(straight) path
 A few deflected nearly
back towards alpha
particle source

Rutherford’s Experiment

Rutherford proposed:
 A very small, dense core at the center of
the atom
 This dense core was called the “nucleus”
 It contains most of the mass of the atom
and it has a positive charge (protons)
 Most of an atom is empty space filled with
electrons
Parts of an Atom

Experimentation in the early 20th century (Thomson
and Rutherford) proved atoms were not indivisible
spheres

Atoms are comprised of smaller particles:
Subatomic particles

More experiments led to the discovery of two more
fundamental subatomic particles: Protons and
neutrons



Electron: Negatively charged (1897)
Proton: Positively charged (1919)
Neutron: No electrical charge (1932)
Nucleus of the Atom (Rutherford Model)

A very dense, small center exists in the
center of the atom called the nucleus


Volume of nucleus is about 1/10 trillionth the volume of
the entire atom
Nucleus is basically the entire mass of the
atom

The protons and neutrons are located in the nucleus

Most of the atom is empty space with fast-moving
electrons
Nucleus of the Atom (Rutherford Model)

The nucleus is the
center (core) of the
atom
 The nucleus
has most of the
mass of the atom
 protons
 neutrons
 The extranuclear
region
 it contains all the
electrons
Extranuclear
region
nucleus
Nucleus of the Atom (Rutherford Model)
 The
nucleus is the core of the atom
Positively charged
Contains most of the mass of the
atom
Within a neutral atom, there are
equal numbers of protons and
electrons, so atom has a net
charge of zero
Mass of Subatomic Particles:
The Proton

The proton:
 Has the same magnitude charge as the electron,
but oppositely charged
 Has a charge of +1
 Weighs about 2000 times an electron
 Is found in the nucleus
 In a neutral atom:
Number of protons = number of electrons

Number of protons = identity of the compound
Mass of Subatomic Particles:
The Neutron
 The
last of the three subatomic
particles to be discovered, also
located in the nucleus
 The mass is about the same as a
proton
 Has no charge (neutral)
 Variable amounts are possible in
atoms of the same element
 This
is the basis for isotopes
Mass of Subatomic Particles

The three subatomic particles
have extremely small masses
 Chemists base the mass of
atoms on the atomic mass
scale
 A relative scale based on the
mass of one carbon atom:
12.00 amu
 One amu is 1/12 the mass of
one carbon atom, so the
approximate mass of one
proton or neutron is 1.00 amu
Atomic Number (Z)
All elements in periodic table arranged
according to the atomic number
 Equal to the number of protons in the
nucleus of an atom

Atomic Number = number of protons in an atom
 Determines
the identity of the atom
 Is also equal to the number of electrons in
the neutral atom
 The top number in each square in the
periodic table
Mass Number (A)

The total number of protons and neutrons in
an atom
Mass Number = number of protons + number of neutrons

Mass number is always a whole number
(no decimals)

An oxygen atom has a mass number of 16
(8 protons and 8 neutrons)
Isotopes and Atomic Mass

All atoms of the same element have the same
atomic number (Z)

The same element can differ in the mass number
(A) due to a different number of neutrons

All Mg atoms have 12 protons, but may have 12,
13, or 14 neutrons
Isotopes and Atomic Mass
 Atoms
that have the same number of
protons and electrons but different
numbers of neutrons are called isotopes
 Since isotopes are atoms of the same
element,
They have the same atomic number
They display the same chemical
properties
Nuclear (Isotopic) Symbols
 A notation
used when necessary to
differentiate between isotopes
A
X
Z
A is the mass number
Z is the atomic number
X is the chemical symbol
Atomic Mass
 A specific
element can have several
mass values if it exists in isotopic forms
 For
example, oxygen atoms can have
any one of three masses but often
treated as if it has one mass
 The
atomic mass of an element is the
mass of the “average atom” of that
element
Atomic Mass

1)
2)
3)
Atomic mass is a
“weighted average
mass” based on:
The number of
isotopes that exist
for the element
The relative mass of
each isotope
The percent
abundance of each
isotope
Example:
Isotopes and Atomic Mass
Complete
Symbol Name
1
1H
19
9F
64
29 Cu
2
1H
the following table:
# Protons #Neutrons #Electrons
Hydrogen
1
0
1
Fluorine
9
10
9
Copper
29
35
29
Hydrogen
1
1
1
end
Remaining
slides (Section 4.6)
will continue with chapter 10
Electron Energy Levels

Electrons possess energy; they are in constant motion in
the large empty space of the atom
 The arrangement of electrons in an atom corresponds to
an electron’s energy
 The electron resides outside the nucleus in one of seven
fixed energy levels
 Energy levels are quantized: Only certain energy values
are allowed
Electron Energy Levels
 Electrons
of similar energy are grouped
into energy levels
 The major energy levels in an atom are
called the principal shells symbolized
by n, the principal quantum number
 As n goes from 1 to 2, 3, 4, etc., the
electron’s energy and distance from the
nucleus increases
 The maximum number of electrons in
an energy level is equal to 2n2
Electron Arrangements
The chemical properties of an element are
determined by the number and arrangement
of electrons about the nucleus
 The electron arrangement (configuration) is
a statement of the number of electrons in
each energy level
 The number of electrons an atom has of
various energies
 The electron arrangements of the first 18
elements can be written by placing electrons
in order of increasing energy

Valence Electrons
 The
electrons that reside in the
highest energy level (n)
 They are the furthest electrons
from the nucleus
 Determine the chemical properties
of an element
 Number at the top of each column
for elements (1A-8A) equals the
number of valence electrons for
each element in that group
Electron-Dot Symbols
 Consists
of an element’s symbol and one
dot for each valence electron placed about
the symbol
 Used only for main group elements
(1A to 8A)
 Main group elements in the same group
have the same number of valence
electrons
 The number of valence electrons is the
same as the group number
Electron-Dot Symbols
Electron-Dot Symbols
 Write
the symbol of the element
 Determine the number of valence
electrons by the group number
 Use dot or X to represent an electron
Group Number
1
Li•
2
3
Be•
•
Li•
4
5
6
7
8
•
••
••
••
••
•B• •C•
•
•
Li+1
•N•
•
••
:F:
•
•O:
•
:F:
•• -1
[:F:]
••
•
:Ne:
••
Electron-dot symbols
Examples
 Determine
the number of valence
electrons for Ba, As, and Br.
 Write the electron-dot symbol for each
of these elements
Ba
 Ba: 2 valence electrons
 As:
5 valence electrons
As
 Br:
7 valence electrons
Br
Ionization Energy
Energy needed to remove one electron from
an atom in the gas state
 It measures how tightly an atom holds its
electrons
 The lower the ionization energy, the easier it
is to remove the electron
 Metals have low ionization energies
 Nonmetals have high ionization energies
 Ionization Energy decreases down the
group
 Ionization Energy increases across the
period
 Left to right

Ionization Energy
 The
more tightly an electron is held,
the higher the ionization energy
 The outermost electrons are the
easiest to remove; as the energy level
increases the farther the electron is
from the nucleus
 As n gets larger, removal of an electron
requires less energy
 Helium requires the most energy of any
element due to its full (stable) energy
level
end