The forces that hold atoms together to form molecules and compounds are known as chemical bonds. Understanding the nature of these bonds is fundamental to predicting a substance’s chemical behavior and physical properties. Every compound must be classified by its primary bond type, which dictates whether component atoms share or fully transfer electrons. Ammonia (\(\text{NH}_3\)), a simple yet common compound, presents a perfect example for exploring this classification challenge. The objective is to definitively determine the type of bond that links the nitrogen and hydrogen atoms in this molecule.
How Atoms Form Chemical Bonds
Chemical bonds are broadly categorized into two primary types: ionic and covalent. This distinction is based on how valence electrons are distributed between the bonded atoms. In an ionic bond, the electron is effectively transferred from one atom to another, creating oppositely charged ions. This typically occurs between a metal and a nonmetal. The resulting electrostatic attraction holds the compound together.
Covalent bonding involves the sharing of electrons between two atoms, usually two nonmetals. If the atoms are identical, the sharing is equal, resulting in a nonpolar covalent bond. When two different nonmetal atoms bond, the sharing is often unequal because one atom pulls the electron cloud closer. This unequal sharing creates a polar covalent bond, where electrons spend more time near one atom than the other, though they are not fully transferred.
Using Electronegativity to Classify Bonds
Chemists use electronegativity (\(\text{EN}\)) to precisely measure an atom’s electron-attracting power within a chemical bond. Electronegativity is defined as an atom’s ability to attract a shared pair of electrons toward itself. The Pauling scale is the most commonly used system for measuring this property.
The difference in electronegativity (\(\Delta\text{EN}\)) between two bonded atoms determines the type of bond they form. A small difference indicates a more equal sharing of electrons, while a large difference points toward a complete transfer. A \(\Delta\text{EN}\) of less than \(0.4\) signifies a nonpolar covalent bond. A difference ranging from \(0.4\) up to \(1.7\) is classified as a polar covalent bond. Any difference greater than \(1.7\) suggests the bond is predominantly ionic.
Determining the Bond Type in Ammonia
To classify the bond in ammonia (\(\text{NH}_3\)), we apply the \(\Delta\text{EN}\) calculation to the nitrogen-hydrogen (\(\text{N-H}\)) bond. On the Pauling scale, the electronegativity value for nitrogen is \(3.04\), and for hydrogen, it is \(2.20\).
Calculating the absolute difference yields \(0.84\) (\(3.04 – 2.20 = 0.84\)). This value falls within the established range of \(0.4\) to \(1.7\), which definitively classifies the \(\text{N-H}\) bond as polar covalent. The bond is polar because the nitrogen atom, having the higher electronegativity value, attracts the shared electrons more strongly. This unequal pull means the electron density is concentrated around the nitrogen nucleus, giving it a partial negative charge (\(\delta^{-}\)), while the hydrogen atoms acquire a partial positive charge (\(\delta^{+}\)).
Why Ammonia Behaves Like a Polar Molecule
The polar nature of the individual \(\text{N-H}\) bonds has consequences for the overall behavior of the ammonia molecule. The nitrogen atom in \(\text{NH}_3\) is bonded to three hydrogen atoms and also possesses a non-bonding lone pair of valence electrons. This arrangement of four electron domains—three bonding pairs and one lone pair—results in a tetrahedral electron geometry around the central nitrogen atom.
The actual shape of the molecule, known as the molecular geometry, is trigonal pyramidal, with the nitrogen atom sitting at the apex. Since the three individual \(\text{N-H}\) bond dipoles are directed toward the more electronegative nitrogen atom, and the molecule is asymmetrical, they do not cancel out. This creates a net dipole moment for the entire \(\text{NH}_3\) molecule, making it a polar molecule. The resulting polarity affects physical properties like solubility, making ammonia highly miscible in water, which is also a polar solvent.