The question of whether hydrogen is slightly negative or positive does not have a single answer because the charge depends entirely on the atom it is chemically bonded to. When atoms join to form molecules, they share electrons, and the resulting charge is rarely a full integer charge, such as H+ or H-. Instead, the concept of a partial charge, represented by the Greek letter delta (\(\delta\)) with a plus or minus sign, describes the uneven distribution of electron density within a bond. This partial charge arises when shared electrons gravitate closer to one atom than the other, making that atom electron-rich (\(\delta-\)) and the other electron-poor (\(\delta+\)). Understanding the rules that govern this electron sharing is the key to determining hydrogen’s charge in any given compound.
The Mechanism of Charge: Electronegativity
The mechanism that determines a partial charge relies on a fundamental atomic property called electronegativity. This is a measure of an atom’s tendency to attract a shared pair of electrons toward itself when part of a chemical bond. The Pauling scale quantifies this tendency, and the difference between the electronegativity values of two bonded atoms dictates the magnitude of the resulting partial charges.
Hydrogen has an intermediate electronegativity value of approximately 2.20. This central position means hydrogen can either lose electron density to a more electronegative partner or gain electron density from a less electronegative partner.
When two atoms with different electronegativities form a covalent bond, the electron cloud is pulled toward the atom with the higher value, creating a polar covalent bond. The more electronegative atom develops a partial negative charge (\(\delta-\)), while the less electronegative atom develops a partial positive charge (\(\delta+\)). This unequal sharing explains why hydrogen’s charge can shift from one molecule to the next.
Hydrogen’s Common State: The Positive Partial Charge (\(\delta+\))
Hydrogen commonly exists with a partial positive charge (\(\delta+\)) because it is typically bonded to elements with significantly higher electronegativity, such as oxygen, nitrogen, and fluorine. These elements are highly electron-hungry and pull the shared electron pair strongly away from the hydrogen nucleus.
Water (H2O) provides the most recognizable example of this phenomenon. Oxygen’s electronegativity is much greater than hydrogen’s, causing the electrons in the O-H bonds to be strongly drawn toward the oxygen atom. As a result, the oxygen atom acquires a partial negative charge (\(\delta-\)), and each hydrogen atom is left with a distinct partial positive charge (\(\delta+\)). This polarity gives water its unique solvent properties and high boiling point.
The partial positive charge on hydrogen is also responsible for the formation of hydrogen bonds, a type of strong intermolecular attraction. This occurs when the electron-poor hydrogen atom on one molecule is electrostatically attracted to a lone pair of electrons on a neighboring highly electronegative atom. These forces are responsible for maintaining the structures of complex biological molecules like proteins and DNA.
Similar partial positive charges are observed when hydrogen bonds to nitrogen (ammonia, NH3) and when it bonds to fluorine (hydrogen fluoride, HF). In all these cases, the non-metal partner’s stronger pull makes the hydrogen atom the electron-deficient, positive pole of the bond.
The Rare Exception: When Hydrogen is Negative (\(\delta-\))
While the partial positive state is most common, hydrogen can acquire a partial negative charge (\(\delta-\)) in a specific, less-frequent scenario. This occurs when hydrogen is bonded to elements that are significantly less electronegative than itself, primarily highly electropositive metals. In these compounds, the metal atom has such a weak hold on its valence electrons that the electron density shifts toward the hydrogen atom.
This situation results in the formation of compounds known as ionic hydrides, where the hydrogen atom formally exists as the hydride ion (H-). For example, in sodium hydride (NaH), sodium is a metal with a very low electronegativity, causing the shared electron density to be pulled toward the hydrogen. Similarly, in complex reducing agents like lithium aluminum hydride (LiAlH4), the hydrogen atoms carry a partial negative charge because they are bonded to the less electronegative metal atoms.
In these metallic hydride compounds, the hydrogen atom behaves chemically as a site of excess electron density, the exact opposite of the electron-poor, slightly positive hydrogen found in water. Although these hydrides are less common in everyday life, they are indispensable reagents in chemical synthesis. This demonstrates that hydrogen is highly versatile and its charge is completely context-dependent.