Chemical bonds are the attractive forces that hold atoms together to form molecules and compounds. The nature of this connection determines nearly all of a substance’s physical and chemical properties, from its melting point to its reactivity. Determining whether the bond within the simple diatomic hydrogen molecule (\(\text{H}_2\)) is ionic or covalent requires understanding how electrons are distributed between the two atoms.
Understanding Chemical Bonds
The fundamental difference between the two main types of chemical bonds lies in the behavior of valence electrons. An ionic bond involves the complete transfer of one or more electrons from one atom to another. This typically occurs between a metal and a nonmetal atom, such as in table salt (\(\text{NaCl}\)). The resulting atoms become oppositely charged ions, which are then held together by electrostatic attraction.
Covalent bonding, by contrast, involves the sharing of valence electrons between two atoms. This type of bond is generally found between two nonmetal atoms that have a similar tendency to attract electrons. By sharing electrons, each atom achieves a more stable configuration, resulting in a cohesive molecule. The equality of that sharing varies widely, which leads to a spectrum of bond types.
The Decisive Factor: Electronegativity
Chemists use the concept of electronegativity to quantify an atom’s ability to attract a shared pair of electrons toward itself within a chemical bond. Developed by Linus Pauling, this scale provides a numerical value for each element, where higher numbers indicate a stronger pull on electrons. Comparing the electronegativity values of two bonding atoms is the primary tool used to classify the type of bond that forms between them.
The difference in electronegativity (\(\Delta\text{EN}\)) between the two atoms determines the bond’s character. If the difference is greater than 1.7 on the Pauling scale, the electron attraction is so unequal that the bond is considered ionic, signifying a near-complete electron transfer. Conversely, if the difference is small, the bond is considered covalent because the electrons are shared relatively equally.
A difference in electronegativity between 0.4 and 1.7 results in a polar covalent bond, where the electrons are shared unequally. The shared electrons spend more time closer to the atom with the higher electronegativity, creating a partial negative charge (\(\delta-\)) on that atom and a partial positive charge (\(\delta+\)) on the other. If the difference is less than 0.4, the bond is classified as nonpolar covalent, indicating a nearly equal sharing.
Analyzing Hydrogen’s Bond
To determine the nature of the bond in the \(\text{H}_2\) molecule, the electronegativity of a single hydrogen atom must be considered. On the Pauling scale, the value for hydrogen is approximately 2.20. Since the \(\text{H}_2\) molecule is formed by two identical hydrogen atoms, the calculation is straightforward.
The calculation of the electronegativity difference (\(\Delta\text{EN}\)) is \(2.20 – 2.20 = 0\). Since the difference is zero, the electron pair is shared perfectly equally between the two hydrogen nuclei. This places the bond firmly in the category of nonpolar covalent. The \(\text{H}_2\) molecule is a classic example of a pure covalent bond, meaning there is no partial charge separation or polarity. Therefore, the bond in \(\text{H}_2\) is definitively covalent, not ionic.