The periodic table
History of the periodic table
Chemists have known about the properties of many elements for centuries, even millennia, but it was not until the 19th century that these elements were put into a logical order. See image 1 and image 2.
Two chemists developed similar classification systems for elements around the same time - Russian Dmitri Mendeleev (1834-1907) and German Julius Meyer (1830-1895). Because Mendeleev's 1869 version of the table preceded Meyer's by a year, the Russian is generally accepted as the 'father' of the periodic table.
Mendeleev arranged known elements in order of their atomic number, the number of protons, increasing from top to bottom, left to right, starting with hydrogen. In addition to ordering elements this way, Mendeleev also noticed that some elements had similar characteristics to others and started grouping these in columns. He found that the elements seemed to fit into regular intervals, or periods, of eight, hence the name periodic table.
Chemists had not discovered all the elements when Mendeleev created the table, but he had the foresight to leave spaces for the discovery of elements that fit the properties of others in its group. Since its creation, chemists have adjusted the table to fit elements that have since been discovered or synthetically developed, but the periodic table as it is known today has changed little from Mendeleev's original format. See image 3.
Groups on the periodic table
The rows on the table are periods and we refer to the columns as groups. Each element possesses similar characteristics to other elements within its group, because of their electron configuration or, more specifically, to the number of electrons they have on their outermost electron shell.
Group I is made up of alkali metals, which are silvery, soft and highly reactive.
Group II contains the alkaline earth metals, which are harder and denser than the alkalis. They are all shiny and silver/white in colour.
The ten columns between group II and III, also known as group IIb, contain a number of elements called transition metals. The transition metals are high-density metals, usually white and shiny, which have high melting points.
Group III sees the start of the metalloids or semi-metals, which possess some characteristics of metals and some characteristics of non-metals. For example, boron is a non-metallic grey powder while the other elements in group III are soft, silvery metals.
Groups IV to VI contain elements of mixed properties including metals and non-metals and elements of different states (solid/gas) at room temperature. Further examination provides more information about the properties of each element.
Group VII belongs to the halogens, which are reactive elements prone to oxidisation (combine with oxygen). The word 'halogen' comes from the Greek term 'to produce salt'.
Group VIII is home to the noble gases; all the elements in this group are gases at room temperature. The noble gases are stable, monatomic (single atom) molecules.
Hydrogen is, once again, the exception because of its single electron. In some chemical reactions it behaves like a group I element, in other cases it acts like a group VII element.
Elements 57 to 71 and 89 to 103 cannot fit into the space they would most logically occupy. They are usually presented separately, below the main bulk of the table, and aligned with the transition metals.
Towards the right of the table is a zigzag, which indicates the presence of metalloids. All the elements to the left of the line, about 80% of the table, are metals, while the rest are non-metals. The rule with metal is that a metallic element becomes more reactive towards the bottom and towards the left of the table, while the opposite is true for the non-metals, excluding the noble gases.
Elemental activity
All atoms have seven electron shells, although simple diagrams depict only occupied shells. Electrons will occupy the inner shells before outer shells, but will not necessarily occupy the inner shells to capacity. See animation 1.
The electron shell closest to the nucleus holds a maximum of two electrons. To represent shell configuration, numbers corresponding with the electrons in each shell are used. Lithium has three electrons, and since the first shell can only hold two electrons, the second holds one, giving the configuration 2,1 to indicate two electrons in the first shell and one in the second. See image 4.
Each electron shell holds a maximum number of electrons. This number is determined by the size of the atom, calculated using the following formula:
Electron shell capacity = 2n 2(where n is the shell number up to seven)
The capacity of the first shell is two, the second shell eight, the third shell 18 and so on.
Electron configuration, the arrangement of electrons, does not directly relate to a shell's capacity. Electron shells fill in stages, the rule being that the outermost shell cannot contain more than eight electrons, regardless of its maximum capacity.
Calcium has 18 electrons, which is deduced from its atomic number of 18. The shell structure of calcium is 2,8,8,2 not 2,8,10 even though the capacity of the third shell is 18. See image 5.
The limit of eight electrons in the outer shell relates to the intervals on the periodic table. Elements with eight electrons in their outer shell are considered more stable than other elements because it takes a lot of energy to add an electron or remove an electron. The group with eight electrons in their outer shells is the noble gases, which are the most stable elements.
Transition metals have up to two electrons in their outermost shell. Their second outermost shells are incompletely filled and they fill their other shells in an irregular manner, which means that their behaviour is sometimes hard to predict in chemical reactions.
The group headings indicate how many electrons are in the outer shell of each element, which reveals their 'combining power'. Valency is the 'combining power' of an element and the electrons on the outer shell are valence electrons. Elements in group I, with one electron on their outer shell, are reactive because they are eager to give away their extra electron. The halogens in group VII, with seven electrons on their outer shell, are also highly reactive because they want to gain an extra electron to fill and stabilise their outer shell. More about valency will learnt from the study of reactions.
