Chemical bonds are the fundamental forces that hold atoms together, forming molecules and larger structures. Understanding these bonds is essential for comprehending the properties and behaviors of substances. Covalent bonds play a primary role in the vast diversity of compounds. Identifying them is a key step in predicting how a substance will interact with its environment.
Understanding Covalent Bonds
A covalent bond forms when atoms share electrons to achieve a more stable electron configuration. Atoms often seek to attain a complete outer electron shell, typically with eight electrons, a principle known as the “octet rule.” By sharing valence electrons, atoms can effectively fill their electron shells, mimicking the stability of noble gases. For instance, two hydrogen atoms, each with one electron, can share their electrons to form a hydrogen molecule, where each hydrogen effectively has two electrons, similar to helium.
A single covalent bond involves the sharing of one pair of electrons between two atoms. When atoms share two pairs of electrons, a double bond forms, and sharing three pairs results in a triple bond. These shared electrons create a strong attractive force that binds the atoms together within a molecule.
The Role of Electronegativity in Bonding
The concept of electronegativity is central to understanding how electrons are shared in a bond and, consequently, the type of bond that forms. Electronegativity quantifies an atom’s ability to attract shared electrons towards itself within a chemical bond. The difference in electronegativity values between two bonded atoms provides a scientific measure to classify the bond as pure covalent, polar covalent, or ionic.
When the electronegativity difference between two atoms is very small, typically less than 0.5, the electrons are shared nearly equally, resulting in a nonpolar or pure covalent bond. If there is a moderate difference, generally between 0.5 and 1.7, the electrons are shared unequally, creating a polar covalent bond where the electrons spend more time closer to the more electronegative atom. A large electronegativity difference, usually greater than 1.7 or 2.0, indicates that electrons are essentially transferred from one atom to another, forming an ionic bond. This spectrum of electron sharing highlights that bonding is not strictly one type or another but rather a continuum.
Identifying Bonds by Element Combination
Covalent bonds typically form between two nonmetal atoms. For example, when hydrogen and oxygen, both nonmetals, combine to form water, they do so through covalent bonds.
In contrast, ionic bonds usually form between a metal and a nonmetal. Metals tend to lose electrons, while nonmetals tend to gain them, leading to the formation of oppositely charged ions that are held together by electrostatic attraction. While this general rule is helpful, some exceptions exist, and the bond character can sometimes have both covalent and ionic contributions.
Observable Properties of Covalent Compounds
Covalent compounds, also known as molecular compounds, exhibit several characteristic macroscopic properties that can aid in their identification. These substances often have relatively low melting and boiling points. This is because the strong covalent bonds exist within the molecules, but the forces between individual molecules (intermolecular forces) are much weaker. Only a small amount of energy is needed to overcome these weak intermolecular forces, not to break the stronger covalent bonds themselves, allowing for easier phase changes.
Covalent compounds generally do not conduct electricity in solid or liquid states. This lack of conductivity stems from the absence of free-moving charged particles, such as ions or delocalized electrons, which are necessary for electrical current to flow. Many covalent compounds exist as gases or liquids at room temperature, distinguishing them from many ionic compounds, which are typically solids with high melting points.