Chemical bonding | Main features | Types of chemical bonding

 

Chemical bonding

Chemical bonding is the permanent attraction between atoms, ions, or molecules that enables the formation of chemical compounds. This bond can be derived from attractive electrical energy by splitting electrons into electrically charged ions, such as ionic bonds or covalent bonds. The strength of chemical bonding varies considerably; there are "strong bonds" or "primary bonds" like covalent, ionic, and metallic bonds and "weak bonds" or "secondary bonds" like dipole-dipole interactions, and hydrogen bonds. In general, strong chemical bonding is involved in the electronic sharing or transfer between participating atoms. Atoms of molecules, crystals, metals, and diatomic gases are present in most of the physical environment around us. These are held together by chemical bonding, which indicates the structure and bulk properties of the material.

Chemical bonding has three main features

i. Atoms interact with each other to form aggregates such as molecules, compounds, and crystals because doing so reduces the total energy of the system; That is, aggregates are more stable than isolated atoms.

ii. Energy is required to separate bonded atoms or ions into isolated atoms or ions. For ionic solids, where ions form a three-dimensional array called a lattice, this force is called lattice energy (U), the enthalpy change occurs when a solid ionic compound is converted into gaseous ions. For covalent compounds, this energy is called bond energy, which is the enthalpy change that occurs when the bond given to the gaseous molecule breaks.

iii. Each chemical bonding is characterized by a special optimal interstitial distance called a bond distance (r0).

Types of chemical bonding

Ionic bonding

Ionic bonding is the electrical interaction between a type of atom that has a large difference in electrical dynamics. There is no exact value for separating ionic form from covalent bonds, but an electromotive difference greater than 1.7 is likely to be ionic, and a difference less than 1.7 is likely to be auxiliary. Ionic bonds lead to individual positive and negative ions. Ionic charges usually range from −3e to + 3e. Ionic bonding usually occurs in metallic salts such as sodium chloride. A feature of ionic bonds is that species take the form of ionic crystals, where no ion is uniquely attached to any other ion in a specific directional bond. Instead, each species of ions is surrounded by ions of reverse charge, and each of the reversely charged ions near it is the same for the same type of surrounding atoms. So it is no longer possible to attach an ion to a specific single ionized atom near it. This is not the case with covalent crystals, where the covalent bonds between specific atoms measured by techniques such as X-ray scattering still differ from the short distances between them.

Covalent bond

The covalent bond is a chemical bonding that involves sharing of electronic pairs into atoms. These electron pairs are known as shared pairs or bond pairs, and the stable balance of attractive and hostile energies between atoms, when they share electrons, is known as a covalent bond. For many molecules, the splitting of the electrons allows each atom to achieve the equivalent of a complete valence shell according to a static electronic configuration. In organic chemistry, covalent bonds are much more common than ionic bonds. Covalent bonds include many types of interactions, including σ-bonds, π-bonds, metal-to-metal bonds, agnostic interactions, bent bonds, three-center two-electron bonds, and three-center four-electronic bonds.

Metallic bond

Metallic bonding is a type of chemical bonding that arises from the electron gravitational force between the conduction electron and the positively charged metal ions. It can be identified as the sharing of free electrons into a structure of positively charged ions (cations). Metal bonding accounts for many of the physical properties of metals such as strength, flexibility, heat and electrical resistance and conductivity, opacity, and luster. Even pure substances, metallic bonds are not only chemical bonds of any kind, which can exhibit metals. For example, the primary gallium is composed of parallelly bound atoms in both liquid and solid states - these pairs form a crystal structure with a metal bond between them. Another example of metal-metal covalent bonds is mercurous ions (Hg²⁺).

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The atoms of metals have a strong attractive force. It takes a lot of energy to overcome it. Therefore, metals often have high boiling points, with tungsten (5828 K) being extremely high. A notable exception is the components of the zinc group: Zn, Cd, and Hg. Their electron configurations end in ... ns², which resembles a great gas configuration like helium, spend more time going down in the periodic table because the energy difference with the empty np orbitals is even greater. These metals are therefore relatively volatile and are avoided in ultra-high vacuum systems. Otherwise, the metal bond can be very strong even in molten metals like gallium. Although gallium will melt from the heat of one's hand just above room temperature, its heating point is not too far away. Molten gallium is therefore a much-uninterrupted liquid thanks to its strong metallic bond.

Hydrogen bonding

The hydrogen bond is the primary force of attraction between the hydrogen (H) atoms that binds to one more electron atom or group and another electron atom that carries a lone pair of electrons - the hydrogen bond acceptor (Ac). Such an interactive system is usually referred to as Dn - H ··· Ac, where the solid line refers to a polar covalent bond and the dotted or dashed line refers to hydrogen bonding. The most frequent donor and receptor atoms are the second row of nitrogen (N), oxygen (O), and fluorine (F).

Hydrogen bonds can be intermolecular or intramolecular, the strength of hydrogen bonds can vary from 1 to 40 kcal/mol depending on the bonds of donor and acceptor atoms, their geometry, and the nature of the environment structure. This makes them somewhat stronger than the van der Waals interactions and weaker than the complete covalent or ionic bonds. This type of bond can occur in inorganic molecules like water and in organic molecules like DNA and proteins. Hydrogen bonds such as paper and felted wool are held together and are responsible for sticking separate sheets of paper together after getting wet and subsequently drying.

Pi bond

In chemistry, pi bonds are covalent chemical bonding where two lobes in the upper orbit of an atom overlap two lobes in an orbit over another atom, and this overlap eventually occurs. Each of these atomic orbits carries two bonded nuclei and has a zero electron density in a shared nodal plane. The same bond is the nodal plane for the molecular orbit of the pi bond. Pi bonds can form in double and triple bonds but in most cases not in single bonds.

The Greek letters π in their name refer to the p orbital as the symmetry of the pi bond orbital is equal to the p orbital as seen in the bond axis. One of the most common forms of this type of bond involves p orbitals, although d orbitals involve pi bonds. This subsequent mode forms part of the base for multiple bonding of metal-metal. Pi bonds are generally weaker than sigma bonds. The bond strength of a C-C double bond consisting of a sigma and a pi bond is less than twice that of a C-C single bond, indicating that the stability added by a pi bond is less than that of a sigma bond.

Sigma bond

A single bond is a sigma bond when multiple bonds are formed by sigma bonds together with pi or other bonds. A double bond has a sigma plus a pi bond, and a triple bond has a sigma plus two pi bonds. In chemistry, sigma bonds (σ bonds) are the strongest forms of covalent chemical bonding. These are formed by head-on overlapping between atomic orbitals. Sigma bonding is most defined for diatomic molecules using the languages ​​and tools of symmetry groups. In this formal method, an σ-bond is symmetrical about the rotation of the bond axis. According to this definition, the most common forms of sigma bonds are s + s, p_z + p_z, s + p_z, and d_z² + d_z². Quantum theory also suggests that molecular orbitals (MO) of uniform symmetry are actually mixed or hybridized. The practical consequence of this mixture of diatomic molecules is that wave-related s + s and p_z + p_z molecular orbits are blended. The degree of this hybridization depends on the relative strength of the symmetrical MOs as preferred.