Periodic Table
Dmitri Mendeleev was the first scientist to create a periodic table of the elements similar to
the one we use today.
This table showed that when the elements were ordered by
increasing atomic weight, a pattern
appeared where properties of the elements repeated periodically. This periodic table is a chart
that groups the elements according to their similar properties.
The most important difference between Mendeleev's table and today's table is the
modern table is organized by
increasing atomic number, not increasing atomic weight.
In 1914,
Henry Moseley learned you could experimentally
determine the atomic numbers of elements. Before that, atomic numbers were just the
order of elements based on increasing atomic weight. Once atomic numbers had significance,
the periodic table was reorganized.
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Periods
Rows of elements are called periods. The period number of an element
signifies the highest unexcited energy level for an electron in that element.
The number of elements in a period increases as you move down the periodic table
because there are more sublevels per level as the energy level of the atom increases.
Groups
Columns of elements help define element groups. Elements within a
group share several common properties. Groups are elements have
the same outer electron arrangement. The outer electrons are called
valence electrons. Because they have the same number of valence electrons,
elements in a group share similar chemical properties. The Roman numerals
listed above each group are the usual number of valence electrons. For example,
a group VA element will have 5 valence electrons.
Families
Most elements are metals. There are so many metals, they are divided into groups:
Alkali metals, Alkaline earth metals, and Transition metals. The transition metals
can be divided into smaller groups, such as the Lanthanides and Actinides.
Group 1: Alkali Metals
The Alkali metals are located in
Group IA (first column) of the periodic table.
Sodium and Potassium are examples of these elements. Alkali metals form salts and many other compounds.
These elements are less dense than other metals, form ions with a
+1 charge, and have the
largest
atom sizes of elements in their periods. The alkali metals are highly reactive.
Group 2: Alkaline Earth Metals
The Alkaline earths are located in
Group IIA (second column) of the periodic table.
Calcium and Magnesium are examples of alkaline earths. These metals form many compounds.
They have ions with a
+2 charge. Their atoms are smaller than those of the alkali metals.
Groups 3-12: Transition Metals
The transition elements are located in groups
IB to VIIIB.
Iron and Gold are examples of transition metals. These elements are very hard,
with high melting points and boiling points. The transition metals are good electrical
conductors and are very malleable. They form positively charged ions.
The transition metals include most of the elements, so they can be categorized into smaller groups.
The
Lanthanides and
Actinides are classes of transition elements. Another way to group transition
metals is into triads, which are metals with very similar properties, usually found together.
Metal Triads
The iron triad consists of iron, cobalt, and nickel. Just under iron, cobalt, and nickel
is the palladium triad of ruthenium, rhodium, and palladium, while under them is the platinum
triad of osmium, iridium, and platinum.
Lanthanides
When you look at the periodic table, you'll see there is a block of two rows of elements below
the main body of the chart. The top row has atomic numbers following
lanthanum. These elements
are called the lanthanides. The lanthanides are silvery metals that tarnish easily. They are
relatively soft metals, with high melting and boiling points. The lanthanides react to form many
different compounds. These elements are used in lamps, magnets, lasers, and to improve the properties
of other metals.
Actinides
The actinides are in the row below the lanthanides. Their atomic numbers follow
actinium.
All of the actinides are radioactive, with positively charged ions. They are reactive metals
that form compounds with most nonmetals. The actinides are used in medicines and nuclear devices.
Groups 13-15: Not all Metals
Groups 13-15 include
some metals, some metalloids, and some nonmetals. Why are these groups mixed?
The transition from metal to nonmetal is gradual. Even though these elements aren't similar enough
to have groups contained within single columns, they share some common properties. You can predict
how many electrons are needed to complete an electron shell. The metals in these groups are called
basic metals.
Nonmetals & Metalloids
Elements that don't have the properties of metals are called nonmetals. Some elements have some,
but not all of the properties of the metals. These elements are called metalloids.
Group 17: Halogens
The halogens are located in
Group VIIA of the periodic table. Examples of halogens are
Chlorine and Iodine.
You find these elements in bleaches, disinfectants, and salts. These nonmetals form ions with a
-1 charge.
The physical properties of the halogens vary. The halogens are highly reactive.
Group 18: Noble Gases
The noble gases are located in
Group VIII of the periodic table.
Helium and Neon are examples of noble gases. These elements are used to make lighted signs,
refrigerants, and lasers. The noble gases are not reactive. This is because they have little
tendency to gain or lose electrons.
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There are mainly four properties which are as follow:
Atomic Radius
The atomic radius of an element is half of the distance between the centers
of two atoms of that element that are just touching each other.
There are three types of atomic radius whuch are as follow:
Metallic radius is half the distance between nuclei in a metallic crystal.
Covalent radius is half the distance between like atoms that are bonded together in a molecule.
Van der Waals radius is the effective radius of adjacent atoms which are not
chemically bonded in a solid, but are presumably in "contact".
An example would
be the distance between the iodine atoms of adjacent I2 molecules in crystalline iodine
Ionization Energy
The ionization energy, or ionization potential, is the
energy required
to completely
remove an electron from a gaseous atom or ion. The closer
and more tightly bound an electron is to the nucleus, the more difficult
it will be to remove, and the higher its ionization energy will be.
The first ionization energy is the energy required to remove one electron from
the parent atom. The second ionization energy is the energy required to remove
a second valence electron from the univalent ion to form the divalent ion, and
so on. Successive ionization energies increase.
The
second ionization energy is
always
greater than the
first ionization energy.
Electron Affinity
Electron affinity reflects the
ability of an atom to accept an electron.
It is the energy change that occurs when an electron is added to a gaseous atom.
Atoms with
stronger effective nuclear charge have
greater electron affinity.
The
Group IIA elements, the alkaline earths, have
low electron affinity
values. These elements are relatively stable because they have filled s subshells.
Group VIIA elements,
the halogens, have
high electron affinities because the addition of an electron to an atom results
in a completely filled shell.
Group VIII elements, noble gases, have
electron affinities near zero,
since each atom possesses a stable octet and will not accept an electron readily. Elements of other
groups have low electron affinities.
Electronegativity
Electronegativity is a measure of the
attraction of an atom for the electrons in a chemical bond.
The higher the electronegativity of an atom, the greater its attraction for bonding electrons.
Electronegativity is related to ionization energy. Electrons with low ionization energies have low
electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements
with high ionization energies have high electronegativities due to the strong pull exerted on electrons
by the nucleus. In a group, the electronegativity decreases as atomic number increases, as a result of
increased distance between the valence electron and nucleus (greater atomic radius). An example of an
electropositive (i.e., low electronegativity) element is cesium; an example of a
highly electronegative
element is
fluorine.
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The properties of the elements exhibit
trends. These trends can be
predicted using the periodic table and can be explained and understood
by analyzing the
electron configurations of the elements. Elements tend
to gain or lose valence electrons to achieve stable octet formation.
Stable octets are seen in the inert gases, or noble gases, of Group VIII
of the periodic table. In addition to this activity, there are two other
important trends.
First,
electrons are
added one at a time moving from
left to right across a
period. As this happens, the electrons of the outermost
shell experience increasingly strong nuclear attraction, so the electrons become
closer to the nucleus and more tightly bound to it. Second, moving down a column
in the periodic table, the outermost electrons become less tightly bound to the
nucleus. This happens because the number of filled principal energy levels
(which shield the outermost electrons from attraction to the nucleus)
increases downward within each group.
These trends explain the periodicity
observed in the elemental properties of atomic radius, ionization energy,
electron affinity, and electronegativity.
Atomic Radius
Generally, the atomic radius decreases across a period from left to right and increases
down a given group. The atoms with the largest atomic radii are located in Group I and
at the bottom of groups.
Moving from
left to right across a
period, electrons are added one at a time to the outer energy shell.
Electrons within a shell cannot shield each other from the attraction to protons. Since the number
of protons is also increasing, the effective nuclear charge increases across a period. This causes
the atomic radius to
decrease.
Moving
down a group in the periodic table, the number of electrons and filled electron shells
increases, but the number of valence electrons remains the same. The outermost electrons in a
group are exposed to the same effective nuclear charge, but electrons are found farther from
the nucleus as the number of filled energy shells increases. Therefore, the
atomic radii increase.
Ionization Energy
Ionization energies
increase moving from
left to right across a
period
(decreasing atomic radius). Ionization energy
decreases moving
down a group
(increasing atomic radius). Group I elements have low ionization energies
because the loss of an electron forms a stable octet.
Electron Affinity
Electron Affinity
increase moving from
left to right across a
period (decreasing
atomic radius). Electron Affinity
decreases moving
down a group (increasing atomic
radius).
Electronegativity
Electronegativity
increase moving from
left to right across a
period
(decreasing atomic radius). Electronegativity
decreases moving
down a group
(increasing atomic radius).
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