By Doug Leenhouts | Updated Aug 30, 2022
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When atoms bond, the resulting structure—whether covalent or ionic—determines the compound’s physical and chemical behavior. Water, for example, is a covalent molecule formed by two hydrogen atoms sharing electrons with one oxygen atom. Understanding these bonds is essential for predicting properties such as melting point, solubility, and reactivity.
Covalent bonds form exclusively between nonmetal atoms that possess similar electronegativities. Because each atom shares one or more electron pairs, the resulting molecules are usually liquids or gases at room temperature and have low melting and boiling points. These compounds are often slightly polar, and the shape of the molecule is defined by the arrangement of shared electron pairs.
A key rule of thumb is that when the electronegativity difference between two atoms is less than 1.7, the bond is covalent. The formation of a covalent bond releases energy, making the compound more stable as additional covalent bonds are formed.
Ionic compounds arise when a metal donates one or more electrons to a nonmetal, creating positively and negatively charged ions that attract each other. An electronegativity difference greater than 1.7 typically indicates ionic character. These solids exhibit high melting and boiling points and are highly polar, reflecting the substantial charge separation.
Many organic molecules are covalently bonded, such as methane (CH₄), where a carbon atom shares electrons with four hydrogens. Covalent bonds can also exist between identical atoms—oxygen (O₂), nitrogen (N₂), and chlorine (Cl₂) are all diatomic gases. Such bonds require significant energy to break, underscoring their strength. In the periodic table, bonds between nonmetals and halogens (group 17) are invariably covalent.
Sodium chloride (NaCl), the familiar table salt, exemplifies an ionic compound. It dissolves readily in water because its ionic lattice is easily disrupted by solvent molecules. Other common ionic substances include magnesium oxide (MgO), potassium chloride (KCl), calcium oxide (CaO), and iron(III) oxide (Fe₂O₃). In each case, atoms achieve a noble‑gas configuration by transferring or accepting electrons, resulting in a stable lattice.
