Chemical bonding describes a variety of interactions that hold atoms together in chemical compounds. Chemical bonds are the connections between atoms in a molecule. These bonds include both strong intramolecular interactions, such as covalent and ionic bonds.
They are related to weaker intermolecular forces, such as dipole-dipole interactions, the London dispersion forces, and hydrogen bonding. The weaker forces will be discussed in a later concept. Chemical bonds : This pictures shows examples of chemical bonding using Lewis dot notation.
Hydrogen and carbon are not bonded, while in water there is a single bond between each hydrogen and oxygen. Bonds, especially covalent bonds, are often represented as lines between bonded atoms. Acetylene has a triple bond, a special type of covalent bond that will be discussed later. Chemical bonds are the forces of attraction that tie atoms together. The nature of the interaction between the atoms depends on their relative electronegativity.
Atoms with equal or similar electronegativity form covalent bonds, in which the valence electron density is shared between the two atoms. The electron density resides between the atoms and is attracted to both nuclei. This type of bond forms most frequently between two non- metals. When there is a greater electronegativity difference than between covalently bonded atoms, the pair of atoms usually forms a polar covalent bond.
The electrons are still shared between the atoms, but the electrons are not equally attracted to both elements. As a result, the electrons tend to be found near one particular atom most of the time. Again, polar covalent bonds tend to occur between non-metals. Finally, for atoms with the largest electronegativity differences such as metals bonding with nonmetals , the bonding interaction is called ionic, and the valence electrons are typically represented as being transferred from the metal atom to the nonmetal.
Once the electrons have been transferred to the non-metal, both the metal and the non-metal are considered to be ions. The two oppositely charged ions attract each other to form an ionic compound. Covalent interactions are directional and depend on orbital overlap, while ionic interactions have no particular directionality. Each of these interactions allows the atoms involved to gain eight electrons in their valence shell, satisfying the octet rule and making the atoms more stable.
Metallic bonding Definition: A metallic bond is formed when the valence electrons are not associated with a particular atom or ion, but exist as a "cloud" of electrons around the ion centers. Metallic materials have good electrical and thermal conductivity when compared to materials with covalent or ionic bonding.
A metal such as iron has metallic bonding. Example: In the real and imperfect world, most materials do not have pure metallic, pure covalent, or pure ionic bonding; they may have other types of bonding as well. For example, iron has predominantly metallic bonding, but some covalent bonding also occurs. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the nonmetal accepts those electrons to become a negatively charged anion. Ionic bonds require an electron donor, often a metal, and an electron acceptor, a nonmetal.
Ionic bonding is observed because metals have few electrons in their outer-most orbitals. By losing those electrons, these metals can achieve noble gas configuration and satisfy the octet rule. Similarly, nonmetals that have close to 8 electrons in their valence shells tend to readily accept electrons to achieve noble gas configuration. In ionic bonding, more than 1 electron can be donated or received to satisfy the octet rule. The charges on the anion and cation correspond to the number of electrons donated or received.
In ionic bonds, the net charge of the compound must be zero. This sodium molecule donates the lone electron in its valence orbital in order to achieve octet configuration.
This creates a positively charged cation due to the loss of electron. This chlorine atom receives one electron to achieve its octet configuration, which creates a negatively charged anion. The predicted overall energy of the ionic bonding process, which includes the ionization energy of the metal and electron affinity of the nonmetal, is usually positive, indicating that the reaction is endothermic and unfavorable. However, this reaction is highly favorable because of the electrostatic attraction between the particles.
At the ideal interatomic distance, attraction between these particles releases enough energy to facilitate the reaction. Most ionic compounds tend to dissociate in polar solvents because they are often polar. This phenomenon is due to the opposite charges on each ion. In this example, the sodium atom is donating its 1 valence electron to the chlorine atom. The most stable state for an atom occurs when its valence electron shell is full, so atoms form covalent bonds, sharing their valence electrons, so that they achieve a more stable state by filling their valence electron shell.
Some covalently bounded compounds have a small difference in charge along one direction of the molecule. This difference in charge is called a dipole, and when the covalent bond results in this difference in charge, the bond is called a polar covalent bond. These kinds of bonds occur when the shared electrons are not shared equally between atoms.
If one atom has a higher electronegativity, the electrons will be drawn closer to the nucleus of that atom, resulting in a small net charge around each nucleus of the atoms in the molecule.
If the atoms in the molecule have the same electronegativity for example, if the atoms are the same, as in N 2 , then the shared electrons will not be drawn towards one nucleus more than another, and the bond will be nonpolar.
Similarly, the higher the difference in electronegativity, the more unequal the sharing of electrons is between the nuclei, and the higher the polarity of the bond. A given nonmetal atom can form a single, double, or triple bond with another nonmetal.
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