Download The Periodic Table of the Elements

The Periodic Table
of the Elements
1869~Dmitri Mendeleev
1944~Glenn Seaborg
History of the Periodic Table
• By 1860, more than 60 elements had been
discovered.
• Different chemists used different atomic
masses for the same elements, resulting in
different compositions being proposed for
the SAME compounds. How could one
chemist understand the results of another?
• In September 1860, the 1st International
Congress of Chemists was held to address
these issues. They agreed on standard
values for atomic masses.
Mendeleev’s Periodic Table
• When the Russian chemist Dmitri Mendeleev
heard about the Congress, he decided to
include the new values for atomic masses in a
textbook he was writing.
• His idea was to organize the elements by
their chemical and physical PROPERTIES
as well as their atomic masses.
• Chemical properties included their “combining
capacity” with other elements, especially with
oxygen and chlorine. Physical properties
included melting and boiling points.
Mendeleev’s Periodic Table
• Mendeleev noticed that when the
elements were organized by increasing
atomic mass, certain similarities in their
chemical properties occurred at regular
intervals. Such repeating patterns are
referred to as PERIODIC.
• Mendeleev created a table in which
elements with similar properties were
grouped together—the first Periodic Table
of the Elements.
Predictions for
Undiscovered Elements
• Mendeleev’s procedure
left several empty
spaces in his table. In
1871, he predicted the
existence and properties
of three of these
elements.
• By 1886, these three
elements—Gallium,
Germanium, and
Scandium—had been
Predicting Properties
• Mendeleev used the average of properties of
elements just above and below an unknown
element to predict its properties.
• Germanium was unknown, but the density of
silicon is 2.3 g/cm3 and tin’s density is 7.3 g/cm3.
• What would you estimate germanium’s density
to be?
When germanium was discovered in 1886, its
density was found to be 5.3 g/cm3, about 10%
different than the prediction of 4.8 g/cm3!
Predicting Properties
• Formulas for chemical compounds can also
be predicted. We know that carbon and
oxygen form carbon dioxide (CO2). What
formula would you predict for a compound of
carbon and sulfur?
Since oxygen and sulfur are in the same group
(16), we can predict that the compound would
be carbon disulfide (CS2).
Your turn!
1. Krypton (Kr) wasn’t known in Mendeleev’s time.
Given that the boiling point of argon (Ar) is
-186°C and of xenon (Xe) is -108°C, estimate
the boiling point of Kr.
2. The melting points of potassium and cesium
(Cs) are 337 K and 302 K respectively.
(a) Estimate the melting point of rubidium (Rb).
(b) Do you expect the melting point of sodium
to be higher or lower than that of rubidium?
Explain the evidence you used for your
prediction.
Solutions
1. The estimated BP of krypton would be
halfway between that of argon and xenon,
or -147°C. (The actual value is -153°C.)
2. (a) The estimated MP of rubidium would
be 320 K. (Actual is 312 K.)
(b) I would predict the MP of sodium to
be higher than that of rubidium as MP
decreases as you go down the group,
and sodium is higher than rubidium in the
group.
Mendeleev’s Principle of
Chemical Periodicity
• Mendeleev’s successful predictions
persuaded most chemists to accept his
periodic table; however, not all elements
could be arranged in order of ATOMIC
MASS.
• In 1911(40 years later) Henry Moseley’s
work on atomic spectra led to the discovery
that the elements fit the pattern of increasing
ATOMIC NUMBER better than atomic mass.
The Periodic Law
• The Periodic Law today is known as:
The physical and chemical
properties of the elements are
periodic functions of their
atomic numbers.
In other words, when the elements
are arranged in order of increasing
atomic number, elements with similar
properties appear at regular intervals.
Periodic Variation in Properties
• Horizontal rows are called PERIODS
• Vertical columns are called GROUPS or FAMILIES.
The elements in each group have similar properties.
 ALKALI METALS—1st column; highly reactive—
form ECl chlorides and E2O oxides
 ALKALINE EARTH—2nd column
 TRANSITION METALS—3rd through 12th column
 HALOGEN—17th column (2nd from right); readily
form 1- anions
 NOBLE GASES—18th column (far right column);
very unreactive
Valence Electrons
• VALENCE ELECTRONS are those
electrons of an atom which are in outer
shells and are available to form
chemical bonds.
• For a main group element, a valence
electron can only be in the outermost
electron shell.
• In a transition metal, a valence electron
can also be in an inner shell.
Valence Electrons for Main Period Groups
• For Groups
1, 2, 13
through 18, it
is easy to
identify the
number of
valence
electrons for
each group.
Valence Electrons for Transition Metals
• A valence electron for a transition metal is defined
as an electron that resides outside a noble-gas
core, which can be in an inner or outer shell.
• The energies of electrons in the d sublevels are
comparable to the s sublevel of the next highest
main level (3d and 4s, 4d and 5s), so an element’s
possible valence electrons are any outside the
noble gas core. Also, the further right in each
series, the lower the energy of an electron in a
d subshell and the less such an electron has
the properties of a valence electron.
• So, the number of valence electrons for
transition metals is hard to predict!
Metals, Metalloids, Nonmetals
Metals
Metalloids
Nonmetals
Types of Elements
• As you can see, the Periodic Table is
broadly divided into two main types of
elements, METALS and NONMETALS
with a few metalloids in between.
• Where are metals located on the periodic
table?
• Where are nonmetals located?
Properties of Metals,
Non-metals, and Metalloids
• Metals have LUSTER (are shiny), are
MALLEABLE (capable of being shaped or
stretched without shattering), and conduct
ELECTRICITY well.
• Nonmetals are usually DULL in appearance,
are BRITTLE, and NONCONDUCTING.
• Metalloids have intermediate properties of
both metals and nonmetals. Silicon and
germanium are examples and are used
extensively in computers.
What determines properties?
• Mostly it is the number and arrangement
of the atom’s electrons.
• Also, metal atoms lose electrons more
readily than nonmetal atoms do.
• Stronger attractions among the atoms of
metals going from left to right in the
Periodic Table means their boiling
points increase.
What do you notice now
about the abrupt change in
BP and MP for Period 2 and 3
elements? Metals are dark
blue, metalloids are lavender,
and non-metals are aqua.
Atomic Bonding
• Most elements do not exist in their ‘pure’ form
but combine with other elements to form
substances with properties different from the
elements themselves.
• A compound is a substance that consists of
two or more elements that are chemically
combined in specific proportions.
• Compounds are formed when atoms combine
with others by gaining, losing, or sharing
electrons, a process called chemical
bonding.
Types of Chemical Bonds
• Ionic bonds form between positive and
negative ions. An ion is an atom that has
an electrical charge through transferral
(loss or gain) of electrons.
Types of Chemical Bonds
• Covalent bonds form when atoms share
electrons. A molecule is the smallest part of a
covalent compound that has the properties
of that compound. Water and many
atmospheric gases such as carbon dioxide are
molecules with covalent bonds.
Types of Chemical Bonds
• Metallic bonds form by sharing a sea of
"free" electrons among a lattice of
positively charged ions.
Properties of Bonds
• Ionic compounds are rigid solids with high
melting and boiling points. When solid, they
are poor conductors of electricity, but when
melted are good conductors. Most are Groups 1
and 2 reacting with Groups 16 and 17.
• Covalent compounds have low melting and
boiling points and have poor electrical
conductivity even when melted.
• Metallic compounds are malleable (easily
shaped), ductile (easily drawn into wires), and
are excellent electrical conductors.
PBS Learning Media
Ionic bonding
Covalent bonding
Trends in the Periodic Table
Electronegativity
Electronegativity
Trend in Atomic Radius
• As you go from left to right across a period,
the atomic radius (distance from nucleus to
outermost electron) tends to decrease. As the
positive charge in the nucleus increases, the
electrostatic attraction on the electrons also
increases, pulling them inward. Remember that
as you go across a period, you stay at the same
energy level.
• The atomic radius usually increases while
going down a group due to the addition of a
new energy level (shell)—the outermost
electrons are farther from the nucleus.
Metallic Character Reactivity
• Character = how much like a metal it behaves;
reactivity = how easily it loses (or gains)
electrons to form chemical bonds
• As you go across a period from left to right,
the nuclear charge increases while the
number of energy levels stays the same, so
there is a stronger and stronger attraction for
the electrons. It becomes more and more
difficult to lose electrons, so the reactivity
of metals decreases as you go from left to
right across the periodic table.
Metallic Character Reactivity
• As you go down a group, the nuclear
charge increases but so does the number
of shielding electrons. This means we
have more and more energy levels and
the electrons are further and further away
from the nucleus, and it is easier for those
electrons to come off.
• So going down a group means a
metal’s reactivity INCREASES.
Ionization Energy
The energy needed to remove one
•
or more electrons from a neutral
atom to form a positively charged
ion is a physical property that
influences the chemical behavior of
the atom.
The first ionization
energy of an
element is the
energy needed to
remove the
outermost, or
highest energy,
electron from a
neutral atom in the
gas phase.
Ionization Energy
• Ionization energy tends to increase
across a period because the greater
number of protons (higher nuclear charge)
attracts the orbiting electrons more strongly,
thereby increasing the energy required to
remove one of the electrons.
• Down a group, the ionization energy will
likely decrease since the valence electrons
are farther away from the nucleus and
experience a weaker attraction to the
nucleus' positive charge.
Second, Third, etc. Ionization Energies
• It will ALWAYS take more energy to remove a
second electron from an ion than it does to
remove the first (outermost) electron from a
neutral atom. Similarly for third, fourth, etc. e• Can it take so much energy to that removing an
electron is unlikely? Consider this:
First, Second, Third, and Fourth Ionization Energies
of Sodium, Magnesium, and Aluminum (kJ/mol)
Electron Affinity
• If ionization energy measures the tendency of a
neutral atom to RESIST the loss of an electron,
electron affinity measures the tendency of a
neutral atom to GAIN an electron.
• It requires energy to remove an electron, but
energy is given off when an electron is gained
(except for Be, N, and noble gases which are
very stable and require energy to accept
another electron!).
• Electron affinities are generally smaller than
ionization energies for reasons we’ll skip for
now.
Electron Affinity Trends
• Going across the Periodic Table from left
to right, electron affinities of the main
group elements tend to INCREASE.
• Going down a group, electron affinities
tend to DECREASE (increase bottom to
top).
• The trends are the same as ionization
energies.
Electronegativity
The tendency of an
atom to attract pairs of
electrons towards itself
for bonding.
An atom's
electronegativity is
affected by both its
atomic number and the
distance that its valence
electrons reside from
the charged nucleus.
Electronegativity and Bonding
• The difference in
electronegativity between
two elements determines
the type of bonding that
occurs between the
elements.
• If the difference is large,
the bond will be IONIC.
• If the difference is small,
the bond will be
COVALENT.
Electronegativity
• Moving left to right across a period,
electronegativity increases due to the
stronger attraction for electrons as the
nuclear charge increases.
• Moving down a group, electronegativity
decreases due to the longer distance
between the nucleus and the valence
electron shell, reducing the attraction for
electrons.
Melting and Boiling Point Trends
Down a Group
What is the trend in BP? In MP?
Why do you suppose is the cause for
these trends?
Melting and Boiling Point Trends
Across Periods 2 and 3
Quick check
Discuss these with a partner and answer:
1. Identify which trend in the diagram below
describes atomic radius:
a)
b)
Increases
Increases
Increases
Increases
c)
Increases
Increases
Quick check
2. Identify which trends below in the
diagram identifies ionization energy,
electron affinity, and electronegativity:
Increases
Increases
Increases
Increases
Increases
Increases
• Why don’t boiling point and melting point
follow a simple trend across a period?
• The answer lies in the type of BONDS that
elements form with other elements to
make chemical compounds.
• One of the trends you did note was an
element’s METALLIC CHARACTER
REACTIVITY. While we will learn about a
metal’s reactivity later, now is a good time
to learn about METALS, METALLOIDS
and NON-METALS.