Understanding the Periodic Table

The periodic table is a way of organising the chemical elements according to their atomic number and properties. With over 110 elements currently reconized a logical system for the classification of these elements is required.

Until the seventeenth century, the views of Aristotle held sway over much of scientific thought. Aristotle believed there were only four elements, which formed all other matter. These elements were earth, air, wind and fire.

During the seventeenth century scientists, such as Robert Boyle, put forward the atomic theory. This theory proposed that there were many elements, each made up of minute particles called atoms. With the acceptance of this theory, scientists examined the world around them looking for new elements. Some materials known from ancient times, such as gold, proved to be elements.

By the early 1800’s, with approaching 50 elements recognized, scientists started to observe similarities in the physical and chemical properties of some elements. In 1863, John Newlands, an English chemist, put the 57 elements known at the time into eleven different groups based on their characteristics.

In 1869, Dimitri Mendeleev, a Russian scientist, put all the then known elements in order of their atomic masses. He then grouped them according to their chemical properties in doing so he produced the periodic table of elements. Within this table Mendeleev, recognized that some elements, such as germanium and scandium, were missing. Thinking that they only waited discovery he left gaps for these missing elements, he also made predictions as to their chemical and physical properties. With the discovery of the missing elements, his predictions proved remarkably correct. His original table did not include the noble gases, as these were not discovered until 1894.

The periodic table is divided into eighteen vertical groups and seven horizontal periods. Elements of the left hand side of the table are metals while those on the right hand side are non-metals.

Under the International Union for Pure and Applied Chemistry (IUPAC) guidelines, groups are numbered 1 to 18 and periods from 1 to 7. Older nomenclature for these groups used Roman numerals combined with the letters A and B. In Europe A and B applies to the left and right side of the table while in America they apply to non-transition and transition elements. Some groups have specific names.

* Group 1 – American IA – European IA – alkali metals.

* Group 2 – American IIA – European IIA – alkaline earth metals.

* Group 3 – American IIIB – European IIIA.

* Group 4 – American IVB – European IVA.

* Group 5 – American VB – European VA.

* Group 6 – American VIB – European VIA.

* Group 7 – American VIIB – European VIIA.

* Group 8 – American VIIIB – European VIIIA.

* Group 9 – American VIIIB – European VIIIA.

* Group 10 – American VIIIB – European VIIIA.

* Group 11 – American IB – European IB – coinage metals.

* Group 12 – American IIB – European IIB.

* Group 13 – American IIIA – European IIIB.

* Group 14 – American IVA – European IVB.

* Group 15 – American VA – European VB – pnictogens.

* Group 16 – American VIA – European VIB – chalcogens.

* Group 17 – American VIIA – European VIIB – halogens.

* Group 18 – American VIIIA – European VIIIB – noble gases.

The older systems are still in common use. They can lead to confusion, for example, an element assigned to group VIA may be a chalcogen in group 16 or a transition metal in group 6 depending on which side of the Atlantic it’s on. In addition, group VIIIB (American) or group VIIIA (European) contains all the elements in groups 8, 9 and 10.

The first element within a group can be used to name that group, so group 16 is called the oxygen group and group 15 the nitrogen group. Other terminologies found in the grouping of elements are transuranium elements (refers to the man-made elements with atomic numbers greater than 93) and super heavy elements (man-made elements with atomic numbers greater than 104. Some of the super heavy elements are awaiting official IUPAC recognition. Until it receives such recognition, IUPAC gives the element a holding name. These holding names begin with the syllable “un”.

Each element in the periodic table has a unique atomic number. The atomic number indicates the number of electrons and protons contained within an atom of that element. Hydrogen, atomic number 1, has 1 proton and 1 electron. Ununoctium, atomic number 118, has 118 electrons and 118 protons.

Protons are positively charged particles held within the nucleus of the atom by neutral neutron particles. Electrons are negatively charged particles that orbit the positively charged nucleus. The electrons orbit within electron shells. Elements in period 1 have 1 electron shell, those in period 2 have 2 electron shells and those in period 7 have 7 electron shells. In all but period 1 these shells are divided into sub-shells, denoted by the letters s, p, d and f. S sub-shells can contain up to 2 electrons, p sub-shells 6 electrons, d sub-shells 10 electrons and f sub-shells 14 electrons.

The presence of sub shells allows the division of the periodic table into blocks. The s-block comprises groups 1 and 2 the p-block is made up of groups 13 to 18. The transition or d-block elements form groups 3 to 11. The separate rows of elements, appearing at the bottom of a chart of the periodic table, form the f-block elements, also known as rare earth elements. There are two series of f-block elements, the lanthanides and the actinides (also known as lanthanoids and actinoids).

Elements form compounds with other elements by losing or gaining electrons. The most reactive metals are the alkali metals, which only need to lose one electron to form a stable ion. The most reactive non-metals are halogens, which only have to gain a single electron to from a stable ion. The most stable and inert of the elements are the noble gases, which have all of their sub-shells filled with electrons. – Read also: noble gases properties

Certain properties show a change across each period. This is particularly noticeable in the first three periods.

Melting points of the elements rise from group 1 to group 14 before dropping markedly in group 15. Melting points continue to drop the element with the lowest melting point for each period is found in group 18.

If a graph is drawn for ionization energy values across a period it shows a general rise across the period from a minimum value in group 1 to a maximum value in group 18. There are two slight dips seen in such a graph at groups 3 and 16 in each period.

Atomic radius decreases along each period but increases down each group.

The periodic table as we know it is now an accepted part of the science of chemistry. It appears in nearly every high school chemistry text book and posters of it decorate the walls of many chemistry laboratories.