phebfid katrinah bernhard
Vita Wanberg
Chemical bonding refers to the formation of a chemical bond between two or more atoms, molecules or ions to give rise to a chemical compound. These chemical bonds are what keep the atoms together in the resulting compound.
The attractive force which holds various constituents (atoms, ions, etc.) together and stabilises them by the overall loss of energy is known as chemical bonding. Therefore, it can be understood that chemical compounds are reliant on the strength of the chemical bonds between their constituents; the stronger the bonding between the constituents, the more stable the resulting compound will be.
The opposite also holds true; if the chemical bonding between the constituents is weak, the resulting compound would lack stability and would easily undergo another reaction to give a more stable chemical compound (containing stronger bonds). To find stability, the atoms try to lose their energy.
Whenever matter interacts with another form of matter, a force is exerted on one by the other. When the forces are attractive in nature, energy decreases. When the forces are repulsive in nature, energy increases. The attractive force that binds two atoms together is known as a chemical bond.
Albrecht Kssel and Gilbert Lewis were the first to explain the formation of chemical bonds successfully in the year 1916. They explained chemical bonding on the basis of the inertness of noble gases.
In 1916, Kossel and Lewis succeeded in giving a successful explanation based upon the concept of an electronic configuration of noble gases about why atoms combine to form molecules. Atoms of noble gases have little or no tendency to combine with each other or with atoms of other elements. This means that these atoms must have stable electronic configurations.
Due to the stable configuration, the noble gas atoms neither have any tendency to gain nor lose electrons and, therefore, their combining capacity or valency is zero. They are so inert that they do not even form diatomic molecules and exist as monoatomic gaseous atoms.
The type of chemical bonds formed varies in strength and properties. There are 4 primary types of chemical bonds which are formed by atoms or molecules to yield compounds. These types of chemical bonds include
Ionic bonding is a type of chemical bonding which involves a transfer of electrons from one atom or molecule to another. Here, an atom loses an electron, which is, in turn, gained by another atom. When such an electron transfer takes place, one of the atoms develops a negative charge and is now called the anion.
The other atom develops a positive charge and is called the cation. The ionic bond gains strength from the difference in charge between the two atoms, i.e., the greater the charge disparity between the cation and the anion, the stronger the ionic bond.
A covalent bond indicates the sharing of electrons between atoms. Compounds that contain carbon (also called organic compounds) commonly exhibit this type of chemical bonding. The pair of electrons which are shared by the two atoms now extend around the nuclei of atoms, leading to the creation of a molecule.
Covalent bonds can be either polar or non-polar in nature. In polar covalent chemical bonding, electrons are shared unequally since the more electronegative atom pulls the electron pair closer to itself and away from the less electronegative atom. Water is an example of such a polar molecule.
A difference in charge arises in different areas of the atom due to the uneven spacing of the electrons between the atoms. One end of the molecule tends to be partially positively charged, and the other end tends to be partially negatively charged.
Compared to ionic and covalent bonding, Hydrogen bonding is a weaker form of chemical bonding. It is a type of polar covalent bonding between oxygen and hydrogen, wherein the hydrogen develops a partial positive charge. This implies that the electrons are pulled closer to the more electronegative oxygen atom.
This creates a tendency for the hydrogen to be attracted towards the negative charges of any neighbouring atom. This type of chemical bonding is called a hydrogen bond and is responsible for many of the properties exhibited by water.
The bond formed as a result of strong electrostatic forces of attraction between a positively and negatively charged species is called an electrovalent or ionic bond. The positively and negatively charged ions are aggregated in an ordered arrangement called the crystal lattice, which is stabilised by the energy called the Lattice enthalpy.
Step 1: Calculate the number of electrons required for drawing the structure by adding the valence electrons of the combining atoms. For example, in methane, a CH4 molecule, there are 8 valence electrons (of which 4 belong to carbon while the other 4 to H atoms).
Step 2: For each negative charge, i.e., for anions, we add an electron to the valence electrons, and for each positive charge, i.e., for cations, we subtract one electron from the valence electrons.
Step 3: Using the chemical symbols of the combining atoms and constructing a skeletal structure of the compound, divide the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.
This structure does not complete octet on N, if the remaining two electrons constitute a lone pair on it. Therefore, we have a double bond between one N and one of the two O atoms. The Lewis structure is
During chemical bonding, when the atoms come closer to each other, the attraction takes place between them, and the potential energy of the system keeps on decreasing till a particular distance at which the potential energy is minimum. If the atoms come closer, repulsion starts, and again, the potential energy of the system begins to increase.
At equilibrium distance, the atoms keep on vibrating about their mean position. The equilibrium distance between the centres of the nuclei of the two bonded atoms is called its bond length.
It is expressed in terms of an angstrom (A0) or picometer (pm). It is determined experimentally by x-ray diffraction or electron diffraction method, or spectroscopic method. The bond length in chemical bonding is the sum of the ionic radii in an ionic compound. In a covalent compound, it is the sum of its covalent radii. For a covalent molecule AB, the bond length is given by d = ra + rb
When atoms come close together, energy is released due to the chemical bonding between them. The amount of energy required to break one mole of bonds of a type so as to separate the molecule into individual gaseous atoms is called bond dissociation enthalpy or bond enthalpy. Bond enthalpy is usually expressed in KJ mol-1.
A bond is formed by the overlap of atomic orbitals. The direction of overlap gives the direction of the bond. The angle between the lines representing the direction of the bond, i.e., the orbitals containing the bonding electrons, is called the bond angle.
In Lewis representation, the number of bonds present between two atoms is called the bond order. The greater the bond order, the greater the stability of the bond during chemical bonding, i.e., the greater the bond enthalpy. The greater the bond order, the shorter the bond length.
In (A), the oxygen-oxygen bond on the left is a double bond, and the oxygen-oxygen bond on the right is a single bond. In B, the situation is just the opposite. The experiment shows, however, that the two bonds are identical.
Therefore, neither structure A nor B can be correct. One of the bonding pairs in ozone is spread over the region of all three atoms rather than localised on a particular oxygen-oxygen bond. This delocalised bonding is a type of chemical bonding in which bonding pair of electrons are spread over a number of atoms rather than localised between two.
Structures (A) and (B) are called resonating or canonical structures, and (C) is the resonance hybrid. This phenomenon is called resonance, a situation in which more than one canonical structure can be written for a species. The chemical activity of an atom is determined by the number of electrons in its valence shell. With the help of the concept of chemical bonding, one can define the structure of a compound, which is used in many industries for manufacturing products in which the true structure cannot be written at all.
These forces occur due to a temporary charge imbalance arising in an atom. This imbalance in charge of the atom can induce dipoles in neighbouring atoms. For example, the temporary positive charge on one area of an atom can attract the neighbouring negative charge.
Atoms having eight electrons in their last orbit are stable and have no tendency to react. Atoms having less than eight electrons then react with other atoms to get eight electrons in their outermost orbit and become stable. Atoms having slightly excess than eight electrons may lose them to atoms which are short of eight. Atoms that cannot either lose or gain may share to get octet configuration. Molecules short of octet configuration, even after the reaction, may accept lone pairs of electrons present in other atoms or molecules.
In metals, the outer orbitals of atoms overlap, and so the electrons present in them do not belong to any particular atom but flow over to all atoms, as well and bind them all together (metallic bonding). Atoms that have to lose and gain electrons become ions and are held together by the electrostatic forces of attraction (ionic bond). When atoms equally give and share electrons, the shared electrons become the unifying force between them (covalent bond). Electron-deficient and free lone pair-containing molecules may again and satisfy the octet thirst of the electron-deficient atom. The shared electron bridges the electron-rich atom with the electron-deficient atom (coordinate bond).
Relatively similar energy sub-orbitals may merge and form a new set of the same number of orbitals, having the property of all the contributing orbitals in proportion to their numbers. These orbitals are hybridized orbitals. They are useful in explaining the similarity in bond length, bond angles, structure, shape and magnetic properties of molecules.
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