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Chapter 4
The Periodic Table
Chapter 4 – The Periodic Table
Section 1 - How Are Elements Organized?
Pure elements at room temperature and atmospheric pressure can be
solids, liquids, or gases. Some elements are colorless. Others are colored.
Despite the differences between elements, groups of elements share
certain properties.
Similarly, the elements fluorine, chlorine, bromine, and iodine can
combine with sodium in a 1:1 ratio to form NaF, NaCl, NaBr, and NaI.
These compounds are also white solids that can dissolve in water to form
solutions that conduct electricity. These examples show that even though
each element is different, groups of them have much in common.
Chapter 4 – The Periodic Table
Section 1 - How Are Elements Organized?
In 1865, the English chemist John Newlands arranged the known
elements according to their properties and in order of increasing atomic
mass. He placed the elements in a table.
In 1869, the Russian chemist Dmitri Mendeleev used Newlands’s observation
and other information to produce the first orderly arrangement, or
periodic table, of all 63 elements known at the time. Mendeleev arranged the
elements in order of increasing atomic mass. Mendeleev started a new row
each time he noticed that the chemical properties of the elements repeated.
Chapter 4 – The Periodic Table
Section 1 - How Are Elements Organized?
He placed elements in the new row directly below elements of similar
chemical properties in the preceding row. He arrived at the pattern
shown below.
Two interesting observations can be made about Mendeleev’s table.
First, Mendeleev’s table contains gaps that elements with particular
properties should fill. He predicted the properties of the missing elements.
Chapter 4 – The Periodic Table
Section 1 - How Are Elements Organized?
Second, the elements do not always fit neatly in order of atomic mass.
For example, Mendeleev had to switch the order of tellurium, Te, and
iodine, I, to keep similar elements in the same column. At first, he thought
that their atomic masses were wrong. However, careful research by others
showed that they were correct. Mendeleev could not explain why his
order was not always the same.
About 40 years after Mendeleev published his periodic table, an English
chemist named Henry Moseley found a different physical basis for the
arrangement of elements. Moseley realized that the periodic table should be
arranged by atomic number, not atomic mass. When the elements were arranged
by increasing atomic number, the discrepancies in Mendeleev’s table disappeared.
Chapter 4 – The Periodic Table
Section 1 - How Are Elements Organized?
The periodic law states that the repeating physical and chemical
properties of elements change periodically with their atomic number
To understand why elements with similar properties appear at regular
intervals in the periodic table, you need to examine the electron configurations
of the elements.
Elements in each column of the table have the same number of electrons
in their outer energy level. These electrons are called valence electrons.
A vertical column on the periodic table is called a group or family.
A horizontal row on the periodic table is called a period.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Elements in groups 1, 2, and 13–18 are known as the main group elements.
Four groups within the main-group elements have special names.
These groups are the alkali metals (Group 1), the alkaline-earth metals
(Group 2), the halogens (Group 17), and the noble gases (Group 18).
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
The Alkali Metals Make Up Group 1
Alkali metals are so named because they
are metals that react with water to make alkaline
solutions.
All have a single valence electron
Alkali metals are usually stored in oil to keep them from reacting with
the oxygen and water in the air.
Because of their high reactivity,
Alkali metals are never found
in nature as pure elements but
are found combined with other
elements as compounds.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
The Alkaline-Earth Metals Make Up Group 2
Like the alkali metals, the alkaline-earth metals
are highly reactive, so they are usually found as
compounds rather than as pure elements
All have 2 valence electrons
The alkaline-earth metals are not as reactive as alkali metals, however
they are harder and have higher melting points.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
The Halogens make up Group 17
Halogens have 7 valence electrons
The halogens are the most reactive group
of nonmetal elements
The halogens have a wide range of physical properties. Fluorine and
chlorine are gases at room temperature, but bromine, is a liquid, and
iodine and astatine are solids.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
The Noble Gases make up Group 18
The noble gases were once called
inert gases because they were
thought to be completely unreactive.
The low reactivity of noble gases leads to
some special uses. Helium, a noble gas, is used to fill blimps because
it has a low density and is not flammable.
Hydrogen Is in a Class by Itself
Hydrogen is the most common element in
the universe. It is estimated that
about three out of every four atoms in the
universe are hydrogen. Because
it consists of just one proton and one
electron, hydrogen behaves unlike any
other element.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
The Transition metals make up Groups 3 through 12
Unlike the main-group elements, the transition
metals in each group do not have identical
outer electron configurations.
Generally, the transition metals are less reactive
than the alkali metals and the alkaline-earth metals are.
Lanthanides and Actinides
Part of the last two periods of transition metals
are placed toward the bottom of the periodic table
to keep the table conveniently narrow,
The elements in the first of these rows are called
the Lanthanides because their atomic numbers
follow the element lanthanum.
Likewise, elements in the row below the lanthanides are called
Actinides because they follow actinium.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Summary of properties of selected groups
Group 1 – All have 1 valence electron, and will easily lose this electron
to form 1+ ions
Group 2 – All have 2 valence electrons, and will easily lose both electrons
to form 2+ ions
Group 13 – All have 3 valence electrons, and will easily lose all of them to
Form 3+ ions
Group 16 – all have 6 valence electrons and will easily gain 2 electron to
Form 2- ions
Group 17 – all have 7 valence electrons and will easily gain 1 electron to
Form 1- ions
Group 18 – have a full outer shell with 8 electrons and will not react with
other atom easily.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Transition elements – are able to have multiple ion charges. All
Possible charges are positive due to lose of electrons. Many will
form a colored aqueous solution.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Most elements of the periodic table are called Metals
All metals are excellent conductors of electricity.
Metals are excellent conductors of heat.
Metals are ductile and malleable.
Ductile means that the metal can be squeezed out into a wire.
Malleable means that the metal can be hammered or
rolled into sheets.
Metals tend to lose electrons to complete the octet rule.
All metals are solid except mercury, which is a liquid.
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Most elements of the periodic table near groups
14 -18 are called Nonmetals
All nonmetals are poor conductors of electricity.
Metals are poor conductors of heat.
Solid nonmetals are brittle and break easily
Metals tend to gain electrons to complete the octet rule.
Nonmetals are found in all three states of matter with bromine being the only
liquid.
Metalloids are elements that share properties of both metals and nonmetals
Metalloids are elements B, Si, As, Te, At, Ge, Sb and Po
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Review Questions
At STP which element has a definite shape and volume?
1. Ag
2. Hg
3. Ne
4. Xe
Which element is a member of the halogen family?
1. K
2. B
3. I
4. S
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Review Questions
The metalloids that are included in Group 15 are antimony (Sb) and
1. N
2. P
3. As
4. Bi
A characteristic of most nonmetallic solids is that they are
1. brittle
2. ductile
3. malleable
4. conductors of electricity
More than two-thirds of the elements of the Periodic Table are classified as
1. metalloids
2. metals
3. nonmetals
4. noble gases
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Review Questions
To which group do the alkaline earth metals belong?
1. 1
2. 2
3. 11
4. 12
Which three elements have the most similar chemical properties?
1. Ar, Kr, Br
2. K, Rb, Cs
3. B, C, N
4. O, N, Si
When metals form ions, they tend to do so by
1. losing electrons and forming positive ions
2. losing electrons and forming negative ions
3. gaining electrons and forming positive ions
4. gaining electrons and forming negative ions
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Review Questions
The chemical properties of the elements are periodic functions of their atomic
1. masses
2. weights
3. numbers
4. radii
Which halogen is a solid at STP?
1. fluorine
2. chlorine
3. bromine
4. iodine
Which element is in Group 2 and Period 7 of the Periodic Table?
1. magnesium
2. manganese
3. radium
4. radon
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Review Questions
An atom in the ground state contains 8 valence electrons. This atom is classified
as a
1. metal
2. semimetal
3. noble gas
4. halogen
Which category is composed of elements that have both positive and negative
charges?
1. the alkali metals
2. the transition metals
3. the halogens
4. the alkaline earths
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Review Questions
Which element is so active chemically that it occurs naturally only in compounds?
1. potassium
2. silver
3. copper
4. sulfur
The elements of the Periodic Table are arranged in horizontal rows according to
each successive element's greater
1. atomic mass
2. atomic radius
3. number of protons
4. number of neutrons
Chapter 4 – The Periodic Table
Section 2 - Tour of the Periodic Table
Review Questions
In which list are the elements arranged in order of increasing atomic mass?
1. Cl, K, Ar
2. Fe, Co, Ni
3. Te, I, Xe
4. Ne, F, Na
Which Group of the Periodic Table contains atoms with a stable outer electron
configuration?
1. 1
2. 8
3. 16
4. 18
The element in Period 4 and Group 1 of the Periodic Table would be classified as a
1. metal
2. metalloid
3. nonmetal
4. noble gas
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
The arrangement of the periodic table reveals trends in the properties of
the elements. A trend is a predictable change in a particular direction.
Metallic Trend
All of the elements of the periodic table increase in metallic
characteristics as they get closer to francium.
This will make group 1 the most metallic group and period 7 the
most metallic period.
Nonmetallic Trend
All of the elements of the periodic table increase in nonmetallic
characteristics as they get closer to flourine.
This will make group 17 the most nonmetallic group and period 2 the
most nonmetallic period.
Note: The noble gases are not considered part of the metallic or nonmetallic trends
due to there resistance to chemical reactions.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Ionization Energy
The energy that is supplied to remove an electron is the ionization energy
of the atom.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Ionization Energy Decreases as You Move Down a Group
Each element has more occupied energy levels than the one above it
has. Therefore, the outermost electrons are farthest from the nucleus in
elements near the bottom of a group.
Similarly, as you move down a group, each successive element contains
more electrons in the energy levels between the nucleus and the
outermost electrons. These inner electrons shield the outermost electrons
from the full attractive force of the nucleus. This electron shielding causes
the outermost electrons to be held less tightly to the nucleus.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Ionization Energy Increases as You Move Across a Period
From one element to the next in a period, the number of protons and the
number of electrons increase by one each. A higher nuclear charge more
strongly attracts the outer electrons in the same energy level, but the
electron-shielding effect from inner-level electrons remains the same.
Thus, more energy is required to remove an electron because the
attractive force on them is higher.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Atomic Radius
The exact size of an atom is hard to determine. An atom’s size depends on
the volume occupied by the electrons around the nucleus, and the electrons
do not move in well-defined paths. Rather, the volume the electrons
occupy is thought of as an electron cloud, with no clear-cut edge. In addition,
the physical and chemical state of an atom can change the size of an
electron cloud.
The diagram below shows one way to measure the size of an atom. This
Method involves calculating the bond radius, the length that is half the distance
between the nuclei of two bonded atoms. The bond radius can change
slightly depending on what atoms are involved.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Atomic Radius Increases as You Move Down a Group
As you proceed from one element down to the next in a group, another
principal energy level is filled. The addition of another level of electrons
increases the size, or atomic radius, of an atom.
Atomic Radius Decreases as You Move Across a Period
As you move from left to right across a period, each atom has one more
proton and one more electron than the atom before it has. As a result,
electron shielding does not play a role as you move across a period.
Therefore, as the nuclear charge increases across a period, the effective
nuclear charge acting on the outer electrons also increases. This increasing
nuclear charge pulls the outermost electrons closer and closer to the nucleus
and thus reduces the size of the atom.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Electronegativity
Atoms often bond to one another to form a compound. These bonds can
involve the sharing of valence electrons. Not all atoms in a compound
share electrons equally. Knowing how strongly each atom attracts bonding
electrons can help explain the physical and chemical properties of the
compound.
Linus Pauling, one of America’s most famous chemists, made a scale
of numerical values that reflect how much an atom in a molecule attracts
electrons, called electronegativity values. Chemical bonding that comes
from a sharing of electrons can be thought of as a tug of war. The atom
with the higher electronegativity will pull on the electrons more strongly
than the other atom will.
Fluorine is the element whose atoms most strongly attract shared
electrons in a compound. Pauling arbitrarily gave fluorine an electronegativity
value of 4.0.Values for the other elements were calculated in relation
to this value.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Electronegativity Decreases as You Move Down a Group
Metals tend to have small electronegativity because they would rather
lose an electron than gain electrons to fulfill the octet rule.
As you go down a group the distance between the valence electrons and
the nucleus increase, weakening the “grip” the nucleus has on it valence
electrons due to electron shielding.
Electronegativity Increases as You Move Across a Period
Nonmetals tend to have large electronegativity because they would rather
gain an electron than lose electrons to fulfill the octet rule.
As you go across a period the distance between the valence electrons and
the nucleus decreases, strengthening the “grip” the nucleus has on it valence
electrons.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Periodic Trends in Ionic Size
Recall that atoms form ions by either losing or gaining electrons. Like
atomic size, ionic size has periodic trends. As you proceed down a group,
the outermost electrons in ions are in higher energy levels. Therefore, just
as atomic radius increases as you move down a group, usually the ionic
radius increases as well. These trends hold for both positive and negative ions.
Metals tend to lose one or more electrons and form a positive ion. As
you move across a period, the ionic radii of metal cations tend to decrease
because of the increasing nuclear charge. As you come to the nonmetal
elements in a period, their atoms tend to gain electrons and form negative
ions. The diagram on the next slide shows that as you proceed through
the anions on the right of a period, ionic radii still tend to decrease
because of the anions’ increasing nuclear charge.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Summary of Ionic Radii
Atoms that form anions will increase in ionic radius, due to the
nuclear charge becoming less than the total electron cloud charge.
Atoms that form cations will decrease in ionic radius due to the nuclear
charge becoming greater than the electron cloud charge.
All atoms forming anions or cations will increase ionic radius when
moving down a group because of the addition of a principle energy
level.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Other information about the periodic table.
All elements with an atomic number greater than 83 have isotopes
that are not stable and will undergo nuclear decay. (They are radioactive)
Chemistry Reference tables lists, electronegativity and ionizations energies
and atomic radii.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Which element forms an ion that is larger than its atom?
1. Cl
2. Ca
3. Li
4. Mg
Which atom has a radius larger than the radius of its ion?
1. Cl
2. Ca
3. S
4. Se
Which of the following elements in Period 3 has the greatest metallic character?
1. Ar
2. Si
3. Mg
4. S
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Which period of the Periodic Table contains more metallic elements than
nonmetallic elements?
1. Period 1
2. Period 2
3. Period 3
4. Period 4
As the elements in Period 3 are considered from left to right, they tend to
1. lose electrons more readily and increase in metallic character
2. lose electrons more readily and increase in nonmetallic character
3. gain electrons more readily and increase in metallic character
4. gain electrons more readily and increase in nonmetallic character
Compared to the nonmetals in Period 2, the metals in Period 2 generally have larger
1. ionization energies
2. electronegativities
3. atomic radii
4. atomic numbers
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Which of the following Group 2 elements has the lowest first ionization energy?
1. Be
2. Mg
3. Ca
4. Ba
Which of the following atoms has the greatest tendency to attract electrons?
1. barium
2. beryllium
3. boron
4. bromine
In which shell are the valence electrons of the elements in Period 2 found?
1. 1
2. 2
3. 3
4. 4
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
Which of the following ions has the smallest radius?
1. F2. Cl3. K+
4. Ca2+
The ability of carbon to attract electrons is
1. greater than that of nitrogen, but less than that of oxygen
2. less than that of nitrogen, but greater than that of oxygen
3. greater than that of nitrogen and oxygen
4. less than that of nitrogen and oxygen
Which trends appear as the elements in Period 3 are considered from left to right?
1. Metallic character decreases, and electronegativity decreases.
2. Metallic character decreases, and electronegativity increases.
3. Metallic character increases, and electronegativity decreases.
4. Metallic character increases, and electronegativity increases.
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
As the atoms in Period 3 of the Periodic Table are considered from left to right,
the atoms generally show
1. an increase in radius and an increase in ionization energy
2. an increase in radius and a decrease in ionization energy
3. a decrease in radius and an increase in ionization energy
4. a decrease in radius and a decrease in ionization energy
Which Group 16 element has only unstable isotopes?
1. Po
2. Te
3. Se
4. S
How much energy is required to remove the most loosely bound electron from
a neutral atom of carbon in the gaseous phase?
1. 801 kJ/mol
2. 1086 kJ/mol
3. 1251 kJ/mol
4. 1000 kJ/mol
Chapter 4 – The Periodic Table
Section 3 - Trends in the Periodic Table
THE END
Review
How is periodic table arranged: In order of atomic number not mass as the
original periodic tables were.
Define: Groups Go up and down on periodic table
Periods Go across the periodic table
Alkali Metals Group 1
Alkaline Earth Metals Group 2
Halogens Group 17
Noble Gases Group 18
Transition Elements Groups 3-12
Metals
Nonmetals
Location
Left
Right
Conduct Electricity
Yes
No
Conduct Heat
Yes
No
Malleable
Yes
No
Brittle (if solid)
No
Yes
Tend to ____ Electrons
Lose
Gain
Name the liquid
Mercury
Bromine
Best represents
Francium
Fluorine
Ionization Energy
Electronegativity
Atomic Radii
Ionic Radii
Energy required to remove one valence electron
Willingness to gain an electron
Size of an atom
Size of an atom after it has gained/lost an electron(s)
Group Trend
Ionization
Energy
Decrease due to increase in
principle energy levels
Period Trend
Increase due to increase
in nuclear charge
Electronegativity
Decrease due to larger radii
Increase due to increase
in nuclear charge
Atomic
Radii
Increase due to increase in
principle energy levels
Decrease due to increase
in nuclear charge
Ionic
Radii
Increase due to increase in
principle energy levels. Does
not matter if gaining or losing
electrons
Varies- If gaining electrons
it will increase if losing
electrons it will decrease