A covalent bond is a chemical linkage formed when atoms share one or more pairs of electrons, typically between nonmetal atoms. Unlike ionic bonding, electrons are shared rather than transferred entirely. The attraction between the shared electrons and the positively charged nuclei holds the resulting stable molecule together. This sharing process is the basis for virtually all organic compounds and many inorganic substances, such as water and carbon dioxide.
Establishing the Need for Electron Sharing
Atoms form covalent bonds primarily to achieve chemical stability. They seek the electron configuration of a noble gas, which represents a stable state. This stability typically means having a completely filled outermost shell, known as the valence shell.
For most main-group elements, stability is achieved with eight valence electrons, following the Octet Rule. Since many atoms start with fewer than eight valence electrons, they must interact with other atoms to complete their shell. By sharing electrons, both atoms can simultaneously count the shared pair toward their stable electron count.
Hydrogen is a notable exception to the Octet Rule, needing only two electrons to fill its outer shell (the Duet Rule). This sharing mechanism is the lowest-energy route to stability for nonmetal atoms, as it allows them to satisfy their shell requirements without the energetic cost of completely losing or gaining multiple electrons. For instance, two hydrogen atoms share their single electron to form a stable Hâ‚‚ molecule.
Step-by-Step Guide to Drawing Covalent Structures
Visualizing how atoms share electrons is accomplished using a Lewis Dot Structure. The first step is to calculate the total number of valence electrons available, creating an “electron budget” for the entire structure. Next, identify the central atom, which is typically the least electronegative atom, noting that hydrogen is always a terminal atom.
The third step involves drawing single bonds connecting each surrounding atom to the central atom. Each single bond uses two valence electrons from the calculated budget. Once placed, the remaining electrons are distributed as lone pairs on the terminal atoms to complete their octets.
Any leftover valence electrons are placed on the central atom as lone pairs. The final check ensures the central atom also possesses an octet. If the central atom is short of eight electrons, a lone pair from an outer atom is moved to form a double or triple bond until the octet is satisfied. For example, in carbon dioxide (\(\text{CO}_{2}\)), two double bonds are formed between the central carbon and each oxygen atom to give all three atoms their octets.
Characterizing the Resulting Bond
Once a covalent structure is established, the resulting bond is characterized by its order and polarity. Bond order refers to the number of shared electron pairs between two atoms. A single bond involves one shared pair, a double bond involves two, and a triple bond involves three. Multiple bonds are progressively stronger and shorter than single bonds due to the increased electron sharing.
The nature of electron sharing is defined by electronegativity, which is an atom’s measure of its tendency to attract electrons toward itself within a chemical bond. The difference in electronegativity between the two bonded atoms determines the bond’s polarity. When two identical atoms bond, such as in the \(\text{H}_{2}\) molecule, the electrons are shared equally because their electronegativity difference is zero, resulting in a nonpolar covalent bond.
If the atoms are different, one atom will have a greater attraction for the shared electrons. The atom with higher electronegativity pulls the electron density closer, acquiring a slight negative charge (\(\delta-\)). This unequal sharing results in a polar covalent bond, creating a molecular dipole. A difference in electronegativity roughly between \(0.5\) and \(1.7\) is considered to indicate a polar covalent bond, while a difference less than \(0.5\) is considered nonpolar.