Silicon is the 14th element of the modern periodic table. It is represented by the symbol ‘Si’ and has an atomic number of 14. Silicon is a group 14 element, along with carbon, germanium, tin, and lead. However, while carbon is a nonmetal, silicon and germanium are considered as metalloids and are very useful in semiconductors. The atomic structure of silicon consists of 14 protons and 14 electrons. The electron configuration of silicon is given by [Ne] 3s23p2. Some more general properties and uses of silicon are discussed in this article.
Because of the high chemical affinity of silicon towards oxygen, it was first isolated into its pure form in the year 1823 by the Swedish chemist Jons Jakob Berzelius. The melting point of silicon is 1414 degrees Celsius whereas the boiling point of silicon is 3265 degrees Celsius. These melting and boiling points are the second highest values among all other metalloids and nonmetals in the modern periodic table. The only nonmetal or metalloid that surpasses the melting and boiling points of silicon is boron.
In the entire universe, the element silicon is the eight most abundant element by mass. However, silicon in its pure only occurs extremely rarely on the Earth’s crust. It is present in the form of dust or sand which is made up of silicon dioxide (SiO2) or other silicates. It is interesting to note that over 90% of the crust of the planet Earth is made up of minerals containing silicates.
In the economy of the modern world, the elemental form of silicon has a huge impact. It has numerous uses including in the refining of steel, the casting of aluminum, and in fine chemical industries. Elemental silicon is very useful in semiconductor electronics and is an essential part of integrated circuits, which is a key part of computers, mobile phones, and other electronic devices.
The skeletal structures secreted by various sea sponges and microorganisms including diatoms and radiolaria are made up of silica. Many plant tissues also contain deposited silica. Trace amounts of silicon are required by animals as well since silicon is an essential element in biology.
The four valence electrons of silicon occupy the 3s and the 3p orbitals. Therefore, silicon can complete the required octet configuration to get the stable noble gas electronic configuration of argon by forming sp3 hybridized orbitals. However, silicon exhibits differences from carbon that are clearly evident. For example, the chemistry of silicon has very few analogies with organic chemistry.
At standard conditions for pressure and temperature, silicon appears to be a shiny substance with a metallic luster that is bluish-grey in color. With the increase in temperature, its resistivity decreases. However, pure silicon is an insulator at room temperatures. The crystal structure of silicon is face-centred diamond cubic. It is one of the p-block elements and corresponds to the third period of the modern periodic table. Silicon is generally doped with group 15 elements to form n-type semiconductors, or doped with group 13 elements to form p-type semiconductors.
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