Chemical bonds are the attractive forces that hold atoms together to form molecules and compounds. These connections are established through the interaction of valence electrons, the electrons in the outermost shell of an atom. The nature of the bond is determined by how these electrons behave, ranging from being shared (covalent) to being fully transferred (ionic). This article provides methods for identifying and characterizing a covalent bond, distinguishing it from an ionic bond.
Identifying Bonds Based on Constituent Atoms
The most straightforward method for preliminary identification of a covalent bond is examining the types of elements involved. Covalent bonds typically form exclusively between two or more nonmetallic elements. This elemental pairing serves as a strong initial indicator that electron sharing, rather than transfer, is occurring.
Nonmetals are characterized by high ionization energies and high electron affinities, meaning they resist giving up electrons but have a strong tendency to gain them. When two nonmetals interact, neither atom is powerful enough to strip an electron completely away from the other. Their similar, strong attractions for electrons result in the valence electrons being shared between the atomic nuclei.
Common examples include the bond between carbon and oxygen in carbon dioxide (\(CO_2\)), or hydrogen and chlorine in hydrogen chloride (\(HCl\)). These shared electron pairs create a stable molecular compound. This contrasts sharply with ionic compounds, which are typically formed from a metal and a nonmetal.
Identifying nonmetals on the periodic table—found generally on the right side, excluding the noble gases—allows for a quick classification. While this elemental rule applies to most cases, it functions best as a preliminary screening tool. For more precise identification, a quantitative approach is required.
Quantifying Bond Type Using Electronegativity
While elemental identity offers a quick classification, the true nature of a bond lies on a continuum, which is best quantified using the concept of electronegativity. Electronegativity is a measure of an atom’s ability to attract a shared pair of electrons toward itself within a chemical bond. This property provides a numerical basis for determining how equally the electrons are being shared.
The quantitative method requires calculating the difference in electronegativity (\(\Delta EN\)) between the two bonded atoms, using values established by the Pauling scale. A smaller difference suggests a more equal sharing, while a larger difference indicates that one atom is pulling the electron density significantly more than the other. This calculated value allows for the precise classification of the bond type.
A difference in electronegativity of less than \(0.4\) is classified as a nonpolar covalent bond. In this scenario, the electron density is shared almost perfectly symmetrically, such as in the bond between two identical atoms, like \(O_2\) or \(N_2\). These bonds are the purest form of electron sharing.
As the difference increases, the bond transitions into a polar covalent bond, generally defined by a \(\Delta EN\) ranging from \(0.4\) up to approximately \(1.7\). In a polar covalent bond, the electron pair is shared unequally, creating a partial negative charge (\(\delta^-\)) on the more electronegative atom and a partial positive charge (\(\delta^+\)) on the less electronegative atom. This charge separation gives the bond a measurable dipole moment.
For instance, the bond between hydrogen (electronegativity \(\approx 2.2\)) and oxygen (electronegativity \(\approx 3.4\)) has a \(\Delta EN\) of about \(1.2\), placing it firmly in the polar covalent range, which explains the high polarity of water. The upper limit of \(1.7\) represents the point where the attraction is so unequal that the bond is considered to have approximately \(50\%\) ionic character.
If the calculated \(\Delta EN\) exceeds \(1.7\), the attractive force of the more electronegative atom is so dominant that the electron is essentially transferred, resulting in an ionic bond. Calculating the numerical difference allows a chemist to precisely locate the bond’s position on the spectrum between purely covalent and purely ionic.
Macroscopic Clues: Physical Properties of Covalent Compounds
Observing a compound’s bulk physical characteristics provides a method for inferring the presence of covalent bonds. Unlike the strong electrostatic attractions within ionic lattices, covalent compounds are characterized by weak intermolecular forces (IMFs) holding the individual molecules together. These weak forces dictate the observable properties of the substance.
A key indicator of a covalently-bonded substance is a low melting point and low boiling point. Because the IMFs are easily overcome by small amounts of thermal energy, many covalent compounds exist as gases or liquids at standard room temperature, such as methane or water. Even solid covalent compounds, like sugar, typically melt at temperatures far below the thousands of degrees required to break the strong ionic bonds in salts.
Another distinguishing property is poor electrical conductivity in all states—solid, liquid, or dissolved in water. Covalent compounds do not dissociate into free-moving ions when melted or dissolved; they remain as neutral molecules. The absence of mobile charged particles means there is no mechanism to carry an electrical current through the substance.
Many covalent compounds exhibit high volatility, meaning they easily transition into the gaseous phase. This tendency is directly related to the low energy required to break the weak intermolecular attractions between molecules. These macroscopic observations—low phase transition temperatures and a lack of conductivity—serve as strong evidence that the underlying chemical structure consists of molecular units held together by covalent bonds.