Periodic Table
Introduction
The Periodic Table is a tabular arrangement of chemical elements, organized based on their atomic number, electron configuration, and recurring chemical properties. Elements are listed in order of increasing atomic number, which corresponds to the number of protons in an element's atomic nucleus. The table has rows called periods and columns called groups.
History
The history of the Periodic Table reflects over a century of growth in the understanding of chemical properties, with major contributions made by Antoine-Laurent de Lavoisier, Johann Wolfgang Döbereiner, John Newlands, Julius Lothar Meyer, Dmitri Mendeleev, and Glenn T. Seaborg.
Early History
The earliest attempt to classify elements was in 1789, when Antoine-Laurent de Lavoisier published a list of 33 chemical elements, grouping them into gases, metals, nonmetals, and earths. However, it was not until the 19th century that scientists started to recognize patterns in the properties of elements and began formulating the first versions of the periodic table.
Development of the Periodic Table
In 1829, Johann Wolfgang Döbereiner observed that many of the elements could be grouped into triads based on their chemical properties. Later, in 1864, John Newlands proposed the Law of Octaves, arranging the elements by atomic weight, noting that similar physical and chemical properties recurred every eighth element.
Julius Lothar Meyer and Dmitri Mendeleev, working independently, both proposed their own periodic tables in 1869 and 1871, respectively. Mendeleev's version gained more acceptance due to his accurate prediction of the properties of yet undiscovered elements.
Modern Periodic Table
The modern periodic table was largely developed by Glenn T. Seaborg starting in the 1940s. He proposed a significant change to Mendeleev's table, introducing the actinide series and reconfiguring the layout of the table.
Structure
The Periodic Table is structured into four blocks: the s-block, p-block, d-block, and f-block. Each block is named after the subshell in which the "last" electron resides.
s-block
The s-block comprises the first two groups (alkali metals and alkaline earth metals) as well as hydrogen and helium. Elements in this block have their last electron in the s orbital.
p-block
The p-block encompasses groups 13 to 18, containing a wide variety of elements including all of the noble gases, the halogens, and various metals and metalloids.
d-block
The d-block consists of groups 3 to 12 and contains all of the transition metals. These elements have their penultimate electron in the d orbital.
f-block
The f-block, often offset below the rest of the periodic table, comprises the lanthanides and the actinides. These elements have their antepenultimate electron in the f orbital.
Periodic Trends
Periodic trends are specific patterns observed in the properties of elements in the Periodic Table. These trends exist because of the similar atomic structure of the elements within their respective groups.
Atomic Radius
The atomic radius of elements generally decreases from left to right across a period and increases down a given group. This is due to the increase in the number of protons and electrons across a period, but the additional electrons are at roughly the same distance from the nucleus.
Ionization Energy
Ionization energy generally increases from left to right across a period and decreases down a group. This is because it is easier to remove an electron from an atom that has a more diffuse electron cloud.
Electronegativity
Electronegativity generally increases from left to right across a period and decreases down a group. This is due to the fact that atoms with high electronegativities tend to have small atomic radii and high ionization energies.