A compound is classified as either ionic or molecular based on how its constituent atoms interact to form a chemical bond. An ionic compound involves the complete transfer of one or more valence electrons from one atom to another, resulting in positively and negatively charged ions held together by strong electrostatic attraction. In contrast, a molecular compound, also called a covalent compound, is formed by the mutual sharing of valence electrons between atoms. Understanding this difference is important because the bond type dictates nearly all of a substance’s observable characteristics, including solubility, state at room temperature, and electrical conductivity.
Identifying the Atoms Involved
Classifying a compound begins by examining the elements it contains and their locations on the periodic table. Ionic compounds typically form between a metal and a nonmetal, which sit on opposite sides of the periodic table. Metals readily lose electrons to form positive ions (cations), while nonmetals readily gain electrons to form negative ions (anions).
Molecular compounds are generally formed from the combination of two or more nonmetal atoms. Since nonmetals have a similar, strong tendency to attract electrons, this leads to the sharing of electrons to achieve a stable configuration. Therefore, when a compound contains only nonmetals, it is almost always a molecular compound. This preliminary check provides a quick classification for the majority of simple compounds.
Using Electronegativity to Determine Bond Type
For a more definitive classification, chemists rely on the concept of electronegativity, which is an atom’s measure of its ability to attract a shared pair of electrons in a chemical bond. The difference in electronegativity ($\Delta EN$) between the two bonded atoms quantifies the degree of electron sharing or transfer. A small difference means the electrons are shared relatively equally, while a large difference indicates a highly unequal sharing that approaches a full transfer.
A $\Delta EN$ greater than approximately 1.7 signifies a predominantly ionic bond, where the electron has been effectively transferred. Bonds with an intermediate $\Delta EN$, typically between 0.4 and 1.7, are classified as polar covalent bonds, meaning the electrons are shared unequally but not completely transferred. A $\Delta EN$ of less than 0.4 results in a nonpolar covalent bond, where electrons are shared nearly equally. This quantitative measure is necessary because the metal/nonmetal rule sometimes oversimplifies the true nature of the bond, which exists on a spectrum between purely covalent and purely ionic.
Recognizing Complex and Exceptional Cases
While the metal/nonmetal pairing is a useful starting point, some compounds do not fit this simple structural rule, requiring a deeper look at the bonding. One common complexity involves compounds containing polyatomic ions, which are groups of atoms that are covalently bonded to each other but carry a net electrical charge.
For example, in sodium sulfate ($\text{Na}_2\text{SO}_4$), the sodium atoms form an ionic bond with the sulfate group ($\text{SO}_4^{2-}$), but the sulfur and oxygen atoms within the sulfate ion are held together by covalent bonds. A compound such as ammonium chloride ($\text{NH}_4\text{Cl}$) is considered an ionic compound because the positive ammonium ion ($\text{NH}_4^+$) and the negative chloride ion ($\text{Cl}^-$) are held together by electrostatic forces. Metalloids, elements like silicon and boron, also create ambiguity, as their $\Delta EN$ values often fall close to the 1.7 boundary, necessitating the use of the quantitative electronegativity rule for accurate classification.
Confirming the Classification through Physical Properties
The bond type dictates the physical properties of a compound, allowing for an observable confirmation of its classification. Ionic compounds are characterized by strong electrostatic attractions that extend throughout a crystal lattice structure, resulting in high melting and boiling points. For instance, sodium chloride melts at over 800 degrees Celsius and is a hard, brittle solid at room temperature.
Molecular compounds, which consist of discrete units with much weaker intermolecular forces between them, exhibit significantly lower melting and boiling points. Many molecular compounds are gases or liquids at room temperature, such as water or carbon dioxide. Ionic compounds conduct electricity when dissolved in water or melted because their charged ions are free to move, whereas most molecular compounds remain poor electrical conductors in any state.