THE PERIODIC TABLE
Chapter 6
Section Overview
• 6.1: Organizing the Elements
• 6.2: Classifying the Elements
• 6.3: Periodic Trends
ORGANIZING THE
ELEMENTS
Section 6.1
Searching for an Organizing Principle
• Chemists used the properties of elements to sort them
into groups.
• J.W. Dobereiner developed a classification system in
which elements were grouped into “triads” or groups of
three with similar properties (ex. Chlorine, bromine, and
iodine form one triad and react easily to metals).
• However, not all of the elements could be so easily
grouped into triads.
• Later a scientist by the name of Mendeleev, arranged the
elements in the periodic table in order of increasing
atomic mass.
The Periodic Law
• In the modern periodic table, elements are arranged in
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order of increasing atomic number, not mass.
Each row is called a period.
Each column is called a group.
Elements in the same period and/or group share
properties.
The periodic law states that when elements are arranged
in order of increasing atomic number, there is a periodic
repetition of their physical and chemical properties.
Metals, Nonmetals, and Metalloids
• Dividing elements into groups is not the only way to
classify them based on their properties.
• The elements can be grouped into three broad classes
based on their general properties.
• The three classes of elements are: metals, nonmetals,
and metalloids.
• Across a period, the properties of elements become less
metallic and more nonmetallic.
Metals, Nonmetals, and Metalloids
• Metals: Good conductors of heat and electric current,
have a high luster (shiny), solid at room temperature,
ductile (drawn into wires), and malleable (hammered into
thin sheets without breaking).
• Nonmetals: Poor conductors of heat and electricity, most
are gases at room temperature, and brittle.
• Metalloids: Properties can behave like both metals and
nonmetals.
CLASSIFYING THE
ELEMENTS
Section 6.2
Squares in the Periodic Table
• The periodic table displays the symbols and names of the
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elements, along with information about the structure of
their atoms.
In the center of the square is the symbol for the element.
Below the symbol is the name of the element.
Above the symbol is the atomic number.
Below the name of the element is the average atomic
mass.
Sometimes, in the top right cornet, the electrons in each
energy level may also be present.
Often the periodic table is color coded to distinguish the
groups of elements.
Squares in the Periodic Table
Electron Configurations in Groups
• Electrons play a key role in determining the properties of
elements.
• Elements can be sorted into noble gases, representative
elements, transition metals, or inner transition metals
based on their electron configurations.
• Noble Gases: The elements in group 8A. The s and p
orbitals in the highest occupied energy level are
completely full, which contributes to their high inactivity.
Helium (He)
1s2
Neon (Ne)
1s22s22p6
Argon (Ar)
1s22s22p63s23p6
Krypton (Kr)
1s22s22p63s23p63d104s24p6
Electron Configurations in Groups
• The Representative Elements: Groups 1A through 7A
display a wide range of physical and chemical properties.
The s and p orbitals of the highest energy level are not
filled. The group number equals the number electrons in
the highest occupied energy level (ex. Ge has 4 and is in
group 4).
Group 1
Lithium (Li)
1s22s1
Group 1
Sodium (Na)
1s22s22p63s1
Group 1
Potassium (K)
1s22s22p63s23p63d104s1
Group 4
Carbon (C)
1s22s22p2
Group 4
Silicon (Si)
1s22s22p63s23p2
Group 4
Germanium (Ge)
1s22s22p63s23p63d104s24p2
Transition Elements
• The two types of transition elements, transition metals
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and inner transition metals, are classified based on their
electron configurations.
Transition metals are the Group B elements that usually
appear in the center of the periodic table.
In atoms of transition metals, the highest occupied s
sublevel and a nearby d sublevel contain electrons, but
classified by the electrons in the d orbitals.
Inner transition metals are below the main body of the
periodic table.
In atoms of inner transition metals the highest occupied s
sublevel and a nearby f sublevel generally contain
electrons, but characterized by f orbitals that have
electrons.
Blocks of Elements
• If you consider both the electron configurations and the
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positions of the elements in the periodic table, another
pattern emerges.
The periodic table can be divided into sections that
correspond to the highest occupied sublevels.
S block: Groups 1A, 2A and Helium.
P block: Groups 3A-8A (exception Helium).
D block: Transition metals
F block: Inner transition metals
The highest occupied energy level is the same number of
the period in which the element is located.
The group number tells how many electrons are in this
energy level.
Blocks of Elements
Blocks of Elements
• Example Problem: Use the “blocks of elements” pattern to
write the electron configuration for nitrogen (atomic
number 7).
Blocks of Elements
• Example Problem: Use the “blocks of elements” pattern to
write the electron configuration for nitrogen (atomic
number 7).
• Solution:
Known
Period = 2
Group = 5
Electrons = 7
1s22p22p3
PERIODIC TRENDS
Section 6.3
Electronegativity decreases
Nuclear charge increases
Shielding increases
Ionic size increases
Ionization energy decreases
Atomic size increases
Periodic Trends Summary
Atomic size decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases
Shielding is constant
Size of cations decreases
Size of anions decreases
Trends in Atomic Size
• Atomic size is expressed as an atomic radius, which is
one half of the distance between the nuclei of two atoms
of the same element when the atoms are joined.
• Atomic radius is measured in picometers (pm).
• In general, atomic size increases from top to bottom
within a group and decreases from left to right across a
period.
• Group trends: As the atomic number increases in a group,
the charge on the nucleus increases and the number of
occupied energy levels increases. The increase in charge
draws electrons towards the nucleus, but the increase in
occupied orbitals shield electrons in the highest level from
the attraction of protons in the nucleus. So, atomic size
increases.
Trends in Atomic Size
• Periodic trends: As you go across a period, each element
has one more proton and one more electron than the
preceding element. The electrons are added to the same
principal energy level and the shielding effect is constant
for all the elements in a period. The increasing nuclear
charge pulls the electrons in the highest level closer to the
nucleus and the atomic size decreases.
Ions
• An ion is an atom or group of atoms that has a positive or
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negative charge.
Positive and negative ions form when electrons are
transferred between atoms.
An ion that has lost an electron and has a positive
charge is called a cation.
An ion that has gained an electron and has a negative
charge is called an anion.
Atoms of metallic elements tend to form ions by losing
electrons from the highest level.
Atoms of nonmetallic elements tend to form ions by
gaining electrons.
Trends in Ionization Energy
• Electrons can move to higher energy levels when atoms
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absorb energy.
The energy required to remove an electron from an atom
is called ionization energy.
This energy is measured in the gaseous state.
The energy required to remove the first electron from an
atom is called the first ionization energy.
First ionization energy tends to decrease from top to
bottom within a group and increases from left to right
across a period.
Trends in Ionization Energy
• Group trends: As the size of an atom increases, nuclear
charge has a smaller effect on the electrons in the highest
orbital. So, less energy is required to remove an electron
from this energy level and the first ionization energy is
lower.
• Period trends: As nuclear charge increases and the
shielding effect remains constant, there is an increase in
the attraction of the nucleus for an electron. So, it takes
more energy to remove an electron from an atom.
Trends in Ionic Size
• During reactions been metals and nonmetals, metals tend
to lose their electrons and nonmetals tend to gain
electrons.
• So, cations are always smaller than the atoms from which
they form and anions are always larger.
• The ionic radii for cations and anions decrease from left to
right across periods and increase from top to bottom
within groups.
Trends in Electronegativity
• Electronegativity is the ability of an atom of an element to
attract electrons when the atom is in a compound.
• In general, electronegativity values decrease from top to
bottom within a group.
• For representative elements, the values tend to increase
from left to right across a period.
• The higher the electronegativity, the more likely that atom
is to “steal” the electrons in a compund.
Electronegativity decreases
Nuclear charge increases
Shielding increases
Ionic size increases
Ionization energy decreases
Atomic size increases
Periodic Trends Summary
Atomic size decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases
Shielding is constant
Size of cations decreases
Size of anions decreases