Main groups and subgroups in the Periodic Table
Table of Contents
- Properties, Electron Configuration, and Atomic Radius
- Main Group Elements
- Subgroup Elements
- Atomic Radius
Properties, Electron Configuration, and Atomic Radius
Here are some facts to get you started: The columns in the Periodic Table of Elements (PSE) are called groups. Historically, the elements were arranged based on their chemical properties. Similarities in the properties of the elements occur periodically. In fact, these similar properties are linked to the valence electron configuration, which is the same across all groups. The effective size of the atoms (the atomic radius) is also important.
Main Group Elements
Main group elements are mostly colorless, diamagnetic (repelled by an electric field), and have paired electrons. The main group number (in a periodic table without subgroups) indicates the number of valence electrons. The valence electrons (the electrons in the outermost shell) of the elements in the main group occupy only s or p orbitals. The remaining shells are completely filled.
The elements on the left in the periodic table have filled s orbitals, and on the right in the PSE, the p orbitals are also filled. Typical metals are found on the left and nonmetals on the right.
The first group is called the alkali metals. Their electron configuration can generally be described as ns1, where n stands for the principal quantum number or period. With the exception of hydrogen, the alkali metals are very reactive due to the unpaired electron in the s orbital. Naturally, they tend to lose their electrons to achieve the noble gas configuration. Reactivity increases from top to bottom of the group. This property is also observed in the alkaline earth metals, but is not as pronounced as in the alkali metals.
The alkaline earth metals belong to the second main group. Their s orbital is filled (ns2), making them less reactive than the alkali metals because two electrons must be removed. Reactivity increases from Mg to Ca (from top to bottom of the group).
The third main group, boron, has the electron configuration ns2(n-1)d10np1. This can be continued for the next groups.
The noble gases have a full p orbital and thus closed shells. Therefore, they are particularly stable and unreactive (chemically inert). They form virtually no compounds.
Subgroup Elements
Subgroup elements are also called transition metals. They are often colored in compounds. Some of them are paramagnetic (attracted by an electric field) and often have unpaired electrons.
The last electron to be filled occupies the d orbital, as in Sc ([Ar]4s2 3d1) or Zn ([Ar]4s2 3d10). The remaining shells are completely filled.
The different oxidation states of manganese exhibit different colors. In one type of experiment, the deep violet potassium permanganate (KMnO4) is poured into a cuvette. A small amount of sodium formate is carefully added using a pipette. During the reaction that takes place, the color in the lower part of the cuvette changes to green after a few minutes. Manganese is thereby reduced, and a very thin blue layer lies between the violet and green layers. A little more sulfuric acid is then added. A pink layer forms at the bottom of the cuvette, which transitions into the green layer via a brown zone.
Lanthanides are also metals and begin in the PSE at La (lanthanum), continue through Ce (cerium), and end at Lu (lutetium). The last electron added is in the 4f orbital.
Actinides are also metals, but radioactive. The last electron added occupies the 5f orbital.
Atomic Radius
Since isolated atoms are essentially infinitely large (because the electrons have enough space to move freely), the atomic radii of elements can only be determined in compounds. The distance between the two atomic nuclei of the atoms in a compound is measured and divided by two (for identical atoms). These radii are called covalent radii.
The atomic radii of the elements increase from top to bottom in the PSE groups, as there are more shells surrounding the nucleus (the orbitals become larger with increasing principal quantum number). From left to right, the atomic radii decrease within a period, as the increasing nuclear charge acts on electrons in the same shells, attracting them more strongly.
The cation of an element is always smaller than the uncharged atom because the positive charge of the nucleus outweighs the negative charge of the electrons, and the electrons are therefore more strongly attracted to the nucleus.
The anion of an element is larger than the uncharged atom because the electrons are less strongly attracted to the nucleus. This means that a cation has the smallest atomic radius, followed by the element and then a singly charged anion. The largest atomic radius is therefore a doubly charged anion, and so on.
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