Periodic Table of Elements

Organization of Periodic Table Periodic Properties of Elements Periodic Trends "Must SEE this" "Interactive Periodic Table"

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.


Organization of Periodic Table


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.


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.


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.


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.


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.


Periodic Properties of Elements

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 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.


Periodic Trends

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 increase moving from left to right across a period (decreasing atomic radius). Electronegativity decreases moving down a group (increasing atomic radius).


Points to Remember

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