The periodic table of elements organizes the chemical elements in a table in order of atomic number (the number of protons in an atom of that element – thus, atoms of an element always retain the same number of protons). Its development is credited to Dmitri Mendeleev (1834-1907), who invented the table on the basis of recurring trends in elemental properties. A row is known as a period, and a column is known as a group. The table looks like this:
Valence electrons are the outermost electrons of an atom, which are important in determining how the atom reacts chemically with other atoms. The number of valence electrons of an element is determined by its column in which the element is categorized. With the exception of groups 3–12 (transition metals), the number within the unit’s place identifies how many valence electrons are contained within the elements listed under that particular column.
Elements’ properties on the periodic table exhibit certain trends. All in all, elements tend to form stable octets, helping to explain these trends.
The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean distance from the nucleus to the outermost valence shell. Atomic radius generally decreases across a period and increases down a group (though with some exceptions such as with phosphorus and sulfur).
The ionization energy is the energy required to remove the outermost electron in the atom when the gas atom is isolated in free space and is in its ground electronic state. Ionization energy generally increases across a period and decreases down a group.
Electron affinity is the measure of the energy released when an electron is added to an atom of the element. Most elements retain negative electron affinities (with Group II – alkaline earth metals – and VIII – noble gas – elements being notable exceptions). Electron affinity increases across a period and decreases down a group.
Electronegativity describes the ability of an atom to attract electrons (or electron density) towards itself. An atom’s electronegativity is affected by both its atomic weight and the distance that its valence electrons reside from the charged nucleus. The higher the associated electronegativity number, the more an element or compound attracts electrons towards it. Like ionization energy and electron affinity, electronegativity generally increases across a period and decreases down a group.
Metallic character describes an element’s likeness to metals in terms of its properties. It is often defined as the tendency of an element to lose electrons and form positive ions, a characteristic of many metals. Metallic character generally decreases across a period and increases down a group.
Here is a summary of periodic table trends:
A metal is generally an element that is a good conductor of both electricity and heat and readily forms cations and ionic bonds with non-metals. They often have high electrical conductivity. They also often retain ductility (the extent to which materials can be pulled into thin strings) and malleability (a material’s ability to form a thin sheet). Metals have low ionization energies and electronegativities.
Nonmetals generally have high ionization energies and are located in the upper right of the periodic table, separated from metals by the diagonal line comprising the metalloids.
- The following nonmetals are diatomic: H2, N2, O2, F2, Cl2, Br2 & I2.
- Phosphorus exists as P4; phosphorus oxide can exist either as P4O10 (most likely) or as P4O6.
- Sulphur is normally present as S8.
The amphoteric line separates metalloids from nonmetals in the periodic table. Metalloids lie along this line. They retain both metallic and non-metallic characteristics. For instance, metalloids are often semi-metals.