Chemical bonding holds atoms together to form molecules and compounds. When atoms interact, the valence electrons in their outermost shells are either transferred or shared. This process determines whether the resulting compound is classified as ionic or covalent. Identifying the bond type is crucial because it dictates a substance’s chemical behavior and physical characteristics, such as melting temperature and solubility.
Identifying Bonds Using Atomic Categories
The most straightforward way to predict a bond type is by locating the atoms on the periodic table. Ionic bonds typically form when a metal and a nonmetal combine, involving a complete transfer of electrons. Metals (left side) lose electrons to form positive ions, while nonmetals (right side) gain electrons to form negative ions. The resulting compound is held together by the strong electrostatic attraction between these oppositely charged particles.
In contrast, covalent bonds usually form between two nonmetal atoms, both found on the right side of the periodic table. Since both atoms strongly attract electrons, neither atom completely transfers an electron. Instead, they achieve stability by sharing valence electrons. This sharing results in a molecular structure rather than an extended network of ions.
A common exception involves polyatomic ions, such as sulfate ($\text{SO}_4^{2-}$) or nitrate ($\text{NO}_3^-$). Within the polyatomic ion, atoms are held together by covalent bonds. However, the interaction between the charged polyatomic ion and a counter-ion (often a metal) is an ionic bond. Compounds like sodium sulfate ($\text{Na}_2\text{SO}_4$) thus exhibit both covalent and ionic characteristics.
The Electronegativity Difference Scale
A more precise method for bond classification involves quantifying the atoms’ electron-attracting power, known as electronegativity. This property measures an atom’s ability to attract shared electrons within a chemical bond. The difference in electronegativity ($\Delta$EN) between two bonded atoms provides a numerical measure of how equally the electrons are shared.
To apply this method, one calculates the absolute difference between the electronegativity values of the two atoms involved. The resulting $\Delta$EN value guides the classification of the bond type. A very small difference, typically less than $0.5$, suggests the electrons are shared almost equally, resulting in a nonpolar covalent bond. Diatomic molecules like oxygen gas ($\text{O}_2$) exemplify this equal sharing.
As the $\Delta$EN increases to a range between $0.5$ and $1.7$, the bond is classified as polar covalent. Here, electron sharing is unequal, causing shared electrons to spend more time near the more electronegative atom. This creates a slight negative charge on that atom and a slight positive charge on the other. Water ($\text{H}_2\text{O}$) contains highly polar covalent oxygen-hydrogen bonds, which accounts for many of its unique properties.
When the electronegativity difference is greater than $1.7$, one atom’s attraction is strong enough to pull the electron completely away. This difference leads to the formation of distinct positive and negative ions, characterizing an ionic bond. The $1.7$ value is not a rigid cutoff, as few bonds are purely $100\%$ ionic or covalent. Most chemical bonds exist on a continuum, but the calculated $\Delta$EN provides a quantitative basis for classification.
Practical Tests Based on Physical Properties
Beyond theoretical calculation, the type of chemical bond can be inferred from a compound’s observable physical properties. The difference between the strong electrostatic forces of ionic bonds and the weaker intermolecular forces of covalent bonds leads to distinct macroscopic behaviors. Testing these properties helps categorize an unknown substance.
One telling characteristic is the melting and boiling point. Ionic compounds, such as table salt, exist as rigid crystal lattices where every ion is strongly attracted to its neighbors. Breaking these numerous strong bonds requires a large amount of energy, resulting in characteristically high melting points, often well above $400^\circ\text{C}$. In contrast, covalent compounds, like sugar, are composed of discrete molecules held together by much weaker forces, meaning they melt and boil at significantly lower temperatures.
Electrical conductivity provides another clear test for bond type. Ionic compounds do not conduct electricity in their solid state because the charged ions are locked in place within the crystal lattice. However, when melted or dissolved in water, the ions become mobile and can move freely to carry an electrical current. Covalent compounds, made of neutral molecules, generally do not produce mobile charged particles and do not conduct electricity in any state.
Finally, solubility patterns often reflect the nature of the bond. Many ionic compounds dissolve readily in polar solvents like water because the charged ions are easily separated by the polar water molecules. Covalent compounds, particularly nonpolar ones such as oils, tend to dissolve better in nonpolar solvents like hexane or carbon tetrachloride. This follows the principle that “like dissolves like,” where materials with similar bonding characteristics mix well.