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.
These atomic properties help describe the macroscopic properties of compounds. For example, smaller covalent compounds that are held together by weaker bonds are frequently soft and malleable. On the other hand, longer-range covalent interactions can be quite strong, making their compounds very durable. Ionic compounds, though composed of strong bonding interactions, tend to form brittle crystalline lattices. Ionic bonds are a subset of chemical bonds that result from the transfer of valence electrons, typically between a metal and a nonmetal.
Ionic bonds are a class of chemical bonds that result from the exchange of one or more valence electrons from one atom, typically a metal, to another, typically a nonmetal.
This electron exchange results in an electrostatic attraction between the two atoms called an ionic bond. An atom that loses one or more valence electrons to become a positively charged ion is known as a cation, while an atom that gains electrons and becomes negatively charged is known as an anion. This exchange of valence electrons allows ions to achieve electron configurations that mimic those of the noble gases, satisfying the octet rule.
The octet rule states that an atom is most stable when there are eight electrons in its valence shell. Atoms with less than eight electrons tend to satisfy the duet rule, having two electrons in their valence shell. By satisfying the duet rule or the octet rule, ions are more stable.
An anion is indicated by a negative superscript charge - something to the right of the atom. Similarly, if a chlorine atom gains an extra electron, it becomes the chloride ion, Cl —. Both ions form because the ion is more stable than the atom due to the octet rule. Once the oppositely charged ions form, they are attracted by their positive and negative charges and form an ionic compound.
Ionic bonds are also formed when there is a large electronegativity difference between two atoms. This difference causes an unequal sharing of electrons such that one atom completely loses one or more electrons and the other atom gains one or more electrons, such as in the creation of an ionic bond between a metal atom sodium and a nonmetal fluorine.
Formation of sodium fluoride : The transfer of electrons and subsequent attraction of oppositely charged ions. To determine the chemical formulas of ionic compounds, the following two conditions must be satisfied:. This is because Mg has two valence electrons and it would like to get rid of those two ions to obey the octet rule.
Fluorine has seven valence electrons and usually forms the F — ion because it gains one electron to satisfy the octet rule. Therefore, the formula of the compound is MgF 2. The subscript two indicates that there are two fluorines that are ionically bonded to magnesium. On the macroscopic scale, ionic compounds form crystalline lattice structures that are characterized by high melting and boiling points and good electrical conductivity when melted or solubilized.
Fluorine has seven valence electrons and as such, usually forms the F — ion because it gains one electron to satisfy the octet rule. Covalent bonds are a class of chemical bonds where valence electrons are shared between two atoms, typically two nonmetals.
The formation of a covalent bond allows the nonmetals to obey the octet rule and thus become more stable. For example:. Covalent bonding requires a specific orientation between atoms in order to achieve the overlap between bonding orbitals. Sigma bonds are the strongest type of covalent interaction and are formed via the overlap of atomic orbitals along the orbital axis. The overlapped orbitals allow the shared electrons to move freely between atoms.
Pi bonds are a weaker type of covalent interactions and result from the overlap of two lobes of the interacting atomic orbitals above and below the orbital axis.
A chemical bond is a lasting attraction between atoms that enables the formation of chemical compounds and may result from the electrostatic force of attraction between atoms with opposite charges, or through the sharing of electrons as in the covalent bonds. The strength of chemical bonds varies considerably. Various theories regarding chemical bonds have been proposed over the past years, during which our interpretation of the world has also changed.
A particular spatial distribution of electrons in a molecule that is associated with a particular orbital energy. As the name suggests, molecular orbitals are not localized on a single atom but extend over the entire molecule.
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