The Periodic Table
And the Periodic Law
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Chapter 6: The Periodic Table
and Periodic Law
• Section 6.1 Development of the Modern
Periodic Table
Antoine Lavoisier
(1743-1794)
• Composed a list in 1790
of the 23 elements
known at the time.
• The advent of
electricity (used to
break down
compounds into their
component elementselectrolysis) and the
development of the
spectrometer, used to
identify elements lead
to the discovery of
many new elements
(70 total by 1870)
Nuclear Magnetic Resonance
(NMR) spectrometer
John Newlands
• Noticed that when elements were arranged
by increasing atomic mass their properties
repeated every eight element.
John Newlands
• This type of pattern is called periodic
because it repeats in a specific manner.
• Newlands called this relationship the Law of
Octaves.
Dmitri Ivanovich Mendeleev
(1834-1907)
• Organized elements increasing atomic
mass
• Noticed that there was a repetition or
periodic pattern in their properties.
Dmitri Mendeleev
• Elements with similar properties were
arranged in columns- this was the first
periodic table
Dmitri Mendeleev
• Mendelev predicted the existence and
properties of undiscovered elements
and left spaces where they should go.
Henry Moseley
• Arranging the elements by mass
resulted in several elements being
placed in groups of elements with
different properties.
Henry Moseley
(1887-1915)
• Discovered that atoms of
each element contained a
unique number of protons
in their nuclei, the number
of protons being equal to
the atom’s atomic number.
Henry Moseley
• By arranging
elements by
increasing atomic
number, the
problems with the
order of the
elements in the
periodic table
were solved.
The Periodic Law
When elements are
arranged by
increasing atomic
number, there is a
periodic repetition
of their chemical
and physical
properties.
The Modern Periodic Table
• Consists of boxes that contain the
elements name, symbol, atomic
number and atomic mass.
The Modern Periodic Table
Groups
• The boxes are arranged in order of
increasing atomic number in a series of
columns called groups or families.
The Modern Periodic Table
Groups
• Each group is numbered 1 through 18.
The Modern Periodic Table
Groups
• Each group is numbered 1 through 8,
followed by the letter A or B.
Periods
• The boxes are arranged in order of
increasing atomic number in a series of
rows called periods.
• Beginning with hydrogen in period 1,
there are 7 periods.
Main Group Elements
• Groups 1,2,and 13-18 are referred to as
the main group (or representative)
elements because they possess a wide
range of chemical and physical
properties.
Transition elements
• Transition elements
• Groups 3-12
Classifying the Elements
• Metals
• All elements to the left side (except H)
of the stair-step line from B to At
Metals
•
•
•
•
•
•
good conductor of heat and electricity
shiny luster
solid at room temperature (except Hg)
Malleable and ductile
high boiling and melting points
High densities
Metals (continued)
• large atomic radius
• low ionization energy
• low electronegativity
Alkali Metals
• The group 1 elements (except for H)
• Most reactive group of metals !!!
Alkaline Earth Metals
• The group 2 elements
• Second most reactive group of metals
Transition Elements
• These are the group B elements (all
are metals)
Transition Metals
• Transition Metals
• Also known as the d-block elements
Inner transition Metals
• Located along the bottom of the
periodic table
• Known as the lanthanide- actinide
series
• Lanthanides are also called rare earth
elements
• Many are radioactive
• Also know as the f- block elements
Non- Metals
• Occupy the upper right side of the
periodic table
• (right of the stair-step line from B to At)
Non-Metals (Characteristics)
•
•
•
•
Generally gasses at room temperature
Poor conductors of heat and electricity
Brittle dull-looking when solids
lower boiling and melting points than
metals (except carbon)
• usually have lower densities than
metals
Non-Metals ( Other Characteristics)
• high electronegativity
• Small atomic radius
• higher ionization
energy than metals
Halogens
• Group 17 elements
• Most highly reactive group of the nonmetals
• Diatomic (Mr. BrINClHOF)
Nobel gases
• Group 18 elements
• Extreamly unreactive
• also called Inert gases
Metalloids
• Elements border the stair-step line from
B to At
• Physical and chemical properties of
both metals and non-metals
Section review 6.1
•
Laviosier,- Composed a list in 1790
of the 23 known elements.
Newlands: Noticed that when elements were arranged by increasing
atomic mass their properties repeated every eight element Law of
Octaves.
Mendeleev : Organized elements increasing atomic mass
Elements with similar properties were arranged in columns- this was the
first periodic table
predicted the existence and properties of undiscovered elements and left
spaces where they should go
Moseley:By arranging elements by increasing atomic number, the
problems with the order of the elements in the periodic table were
solved
.
• Metals- good conductor of heat and electricity, shiny luster,
solid at room temperature (except Hg), Malleable and ductile,
high boiling and melting points, High densities.
• Non-metals- Generally gasses at room temperature , Poor
conductors of heat and electricity, Brittle dull-looking when
solids, lower boiling and melting points than metals (except
carbon), usually have lower densities than metals.
• Metaloids have properties between metals and non metals.
•
•
•
•
A. RE
B. TE
C. TE
D. RE
• A. F and Cl
• B. Be and Mg
• C. Ru and Os
• Ge
Classification of the Elements
By noting an atoms position on the
periodic table, you can determine an its
electron configuration and number of
valence electrons
Valence Electrons
electrons in an atoms highest principle
energy level
• elements have similar chemical
properties because they have the same
number of valence electrons
Valence Electrons and Period
The energy level of an elements valence
electrons indicates the period on the
periodic table in which it is found
• Ex. Li - found in period 2; valence
electrons found in second energy level
• Ex. Gallium (Ga) found in period 4;
valence electrons found in fourth
energy level
Valence electrons and group number
• A main group (representative) element
and the number of valence electrons it
contains are related.
• Group 1 has one valence electron,
Group 2 has 2 valence electrons and so
on. ( Exception He is in group 18, but
only has 2 valence electrons
• Lewis electron dot structures illustrate
this connection.
The s- p- d- and f- block
elements
• The periodic
table is divided
into sections,
or blocks
representing
the atoms
energy sub
levels being
filled with
valence
s- block elements
Consist of groups 1 and 2, H and He
s- block elements
• Valence
electrons occupy
only s orbitals
• group 1 elements
have partially
filled s orbitals
s1
• group 2 elements
have filled s
orbitals s2
s- block elements
• Group 1
elements have
partially filled s
orbitals s1
• Group 2
elements have
filled s orbitals
s2
p- block elements
• After the s sublevel is full, the valence
electrons next occupy the p sublevel
with its three p orbitals.
p- block elements
• The p-block is comprised of groups 13
through 18, with filled or partially filled
p orbitals
p- block elements
• There are no p block elements in the
period 1, the first p-block element B is
in period 2
p- block elements
• Together the
s- and pblocks
comprise the
main group
elements.
p- block elements
• Group 18 are the Nobel gases that are
unreactive with any other elements.
p- block elements
• Nobel gases have their highest PEL
completely full, this is a very stable
electron configuration.
d- block elements
• Largest of the blocks
• Contains the transition metals
d- block elements
• Characterized by a filled outermost s
orbital and filled or partially filled d
orbitals
• d- block spans
10 groups on
the periodic table
f- block elements
• Contains the inner transition metals
(lanthanide-actinide series)
f- block elements
• Have filled or partially filled outermost
s orbital, and filled or partially filled 4f
and 5f orbitals
F- block
• Electrons do not fill their orbitals in a
predictable manner.
F- block
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• There are 7 seven
f orbitals that
hold a maximum
of 14 electrons,
so the f-block
spans 14
columns of the
periodic table.
http://freezeray.com/flashFiles/atomicStructure.htm
• Period 1: s-block elements only
• Periods 2 and 3: s- and p- block
elements
• Periods 4 and 5: s-, p-, and d- block
elements
• Periods 6 and 7: s-, p-, d- and f- block
elements
Section 6.3 Periodic Trends
• As you move across a period or down a
group elements tend to change in a
predictable way, this is known as a
periodic trend.
Atomic Radius
• Defined as half the
distance between adjacent
nuclei in a crystal of the
element
• For diatomic atoms AR is
defined as half the
distance between nuclei of
identical atoms that are
chemically bonded
together.
Trends within periods
• In general there
is a decrease in
atomic radii as
you move leftto-right across a
period
Trends within periods
• In general there
is a decrease in
atomic radii as
you move leftto-right across a
period
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• The result is that the increased nuclear
charge pulls the outermost electrons
closer to the nucleus.
Trends within groups
• Atomic radii generally increase as you
move down a group.
• Electrons are added to higher PELs, so
the outermost orbitals increase the size
of the atom.
• The valence electrons are also farther
from the nucleus
• Largest: Na
• Smallest: S
• Largest: Xe
• Smallest: He
• No, with only this information, you will be unable
to determine the specific groups and periods
That the elements are in; so you can’t apply the
periodic in atomic size to determine which
element has the larger atomic radius.
• No, with only this information, you will be
unable to determine the soecific groups
and periods That the elements are in; so
you can’t apply the periodic in atomic size
to determine which element has the larger
atomic radius.
Ionic Radius
• When atoms lose electrons they form positively
charged ions, they always become smaller.
Ionic Radius
• This is due to the loss of electrons that are
always valence electrons, thus the outer orbital
is empty resulting in a smaller radius.
Ionic Radius
• Atoms can gain or lose one or more
electrons to form ions and gain a net
charge.
Ionic Radius
• Atoms can gain or lose one or more
electrons to form ions and gain a net
charge. An ion is an atom or bonded group
of atoms that has a positive or negative
charge.
Ionic Radius
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• When atoms gain electrons, they have a
negative charge resulting in a larger
radius.
Trends within periods
Ionic Radius
As you move left to
right across the
period the ionic
radius decreases,
then when you get
to group 5A or 6A,
the size of the
larger negative ions
also decreases.
• Trends within
groups
• As you move down
a group, an ion’s
outer electrons are
in a higher PEL
resulting in a
gradual increase in
size.
Ionic radius
• Atoms can gain or lose one or more
electrons to form ions and gain a net
charge. An ion is an atom or bonded group
of atoms that has a positive or negative
charge.
Ionic radius
• When atoms lose electrons they form
positively charged ions, they always
become smaller. This is due to the loss of
electrons that are always valence
electrons, Thus the outer orbital is empty
resulting in a smaller radius.
Ionic radius
When atoms gain electrons, they have a
negative charge resulting in a smaller
radius.
Ionic radius
• Trends within
periods
• As you move left to
right across the
period the ionic radius
decreases, then when
you get to group 5A
or 6A, the size of the
larger negative ions
also decreases.
Ionic radius
Trends within groups
• As you move down a group, an ion’s outer
electrons are in a higher PEL resulting in a
gradual increase in size.
Ionization energy
• To form a positive ion, an electron must be
removed from a neutral atom.
• Ionization energy is defined as the energy
required to remove an electron (from the
gaseous state).
• The energy required to remove the first
electron is called the first ionization energy
(energy required to remove the second
electron, is called the second ionization
energy)
Ionization energy
• A high ionization energy value indicates
that an atom has a strong hold on
electrons.
• IE is a measure of how strongly an
atom’s nucleus will hold onto valence
electrons.
• Atoms with large IE are less likely to form
positive ions than atoms with low IE.
Ionization energy
• Trends within periods
• First ionization energies generally increase
as you move l to r across a period.
Ionization energy
Trends within groups
• First ionization energies generally
decrease as you move down a group.
• With the valence electrons farther away
from the nucleus, less energy is required
to remove them.
Octet Rule
• Atoms tend to gain, lose or share
electrons in order to acquire a full set of
eight valence electron, (noble gas
configuration).
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Octet Rule
• This is a very stable electron configuration.
Elements to the right of the periodic table
tend to gain electrons (form negative ions)
• Elements to the left of the periodic table
tend to lose electrons (form positive ions
to attain this configuration.
Electronegativity
• The relative ability of an atom to attract
elements in a chemical bond.
• Expressed in terms of a numerical value of
3.98 or less.
• Units called Paulings (see Linus Pauling
1901-1994)
Electronegativity
• The relative ability of an atom to attract
elements in a chemical bond.
Electronegativity
• Noble gases have been left out
• F is the most electronegative element
(3.98); Cs and Fr are the least
electronegative (.79 and .7)
• In a chemical bond the atoms with the
greater electronegativity more strongly
attract the bond’s electrons
Electronegativity
Trends within groups and periods
• Electronegativity generally decreases as
you move down a group and increases as
you move left to right across a period.
Electronegativity
Trends within groups and periods
• Elements with the higher electronegativity
are at the upper right and elements with
the lower electronegativity are at the
bottom left of the periodic table.
AR increase down a group as the electrons are added to higher
energy levels and inner core electrons shields the valence electrons
from the increased nuclear charge. AR decrease across a period as
increased nuclear charge coupled with unchanging shielding by inner
core electrons pulls the valence electrons (being added to the same
energy level) closer to the nucleus.
• Largest: antimony (Sb)
Smallest: (nitrogen (N)
a. F
b. Br
c. Br
d. F
• Lithium’s second removed electron is an inner core electron, not a
valence electron. Carbon’s fourth removed electron is still a valence
electron.