Ionic Bond
An ionic bond is formed through the complete transfer of electrons from one atom to another, typically between a metal and a non-metal. For instance, sodium (Na), a metal, loses an electron to become positively charged (Na+). Conversely, chlorine (Cl), a non-metal, gains an electron to become negatively charged (Cl-). This creates ions with opposite charges that are strongly attracted to each other, like the poles of magnets. The resulting compound, such as sodium chloride (NaCl), exhibits a crystal lattice structure, where ions are arranged in a repeating pattern that extends in all three dimensions.
Attributes of ionic bonds include high melting and boiling points, due to the strong forces holding ions together. These compounds also conduct electricity when melted or dissolved in water, as the mobile ions can carry an electric current. However, ionic solids tend to be brittle; any distortion in the lattice can result in like-charged ions being forced close to each other, repelling and causing the structure to fracture.
Covalent Bond
In a covalent bond, atoms share electrons to achieve a full outer shell, resembling a stable noble gas configuration. This occurs primarily between non-metal atoms with similar electronegativities. The shared electrons occupy the space between the atoms and contribute to both atoms' outer shells. We observe this electron sharing in a molecule of carbon dioxide (CO2), where carbon and oxygen, both non-metals, share electrons to form double bonds.
Covalent compounds have distinct properties: they typically have lower melting and boiling points compared to ionic compounds, and they do not conduct electricity, as there are no free-moving charged particles. The shared electron pairs in covalent bonds can be single, double, or triple, with triple bonds being the strongest and shortest. A perfect example is the nitrogen molecule (N2), where two nitrogen atoms share three pairs of electrons, creating an exceedingly strong triple bond.
Electronegativity
Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a chemical bond. The periodic table shows a general trend for electronegativity: it increases as you move from left to right across a period and decreases as you move down a group. Atoms with higher electronegativity, such as oxygen, fluorine, and chlorine, tend to attract electrons more strongly.
Differences in electronegativity between bonding atoms influence the type of bond formed. A large difference usually results in an ionic bond, while a smaller or non-existent difference leads to a covalent bond. When two atoms have exactly the same electronegativity, they share electrons equally, as seen in the N2 molecule; however, when atoms like C and O bond, whose electronegativities are different, the sharing is not equal, leading to poles or partial charges developing within the molecule, also called a polar covalent bond.
Electron Sharing
Electron sharing is the hallmark of covalent bonding and is essential for understanding how molecules form and behave. This sharing allows each atom in the bond to count the shared electrons towards their valence, helping them attain a full outer electron shell. In covalent bonds, the number of shared electron pairs can vary, leading to single, double, or triple bonds, as seen in compounds like methane (CH4), oxygen gas (O2), and nitrogen gas (N2), respectively.
Electron sharing can be equal (nonpolar) when the bonded atoms are identical or have very close electronegativity values, or it can be unequal (polar) if there's a significant difference in electronegativity. Unequal sharing results in partial charges on the atoms, with the more electronegative atom gaining a slight negative charge and the less electronegative atom a slight positive charge. This aspect of electron sharing is critical in the creation of molecules with complex shapes and specific interactions, which are vital for the diverse chemistry of life.
