Hydrogen occupies a unique and transitional position on the periodic table, often behaving in ways that defy simple classification. With only one proton and one electron, this element readily forms chemical bonds to achieve a stable electron configuration. In most compounds encountered in daily life, hydrogen is known to share its single electron, which leaves it with a slightly positive charge. This common behavior leads to the central question of whether hydrogen can ever reverse its role and carry a negative charge. Understanding how charge distributes across molecules helps clarify that hydrogen’s behavior is flexible and determined entirely by its bonding partner.
The Concept of Electronegativity and Bond Polarity
The distribution of electrons in a chemical bond is governed by a measurable atomic property called electronegativity. This property describes an atom’s inherent power to attract a shared pair of electrons toward itself within a bond. To quantify this electron-pulling strength, scientists use scales like the Pauling scale, which assigns a numerical value to each element, typically ranging from about 0.7 to 4.0.
When two different atoms bond together, their unequal electronegativity values determine the nature of the bond. If one atom has a significantly stronger pull than the other, the shared electrons will spend more time orbiting the nucleus of the more attractive atom. This unequal sharing creates an electric dipole, a separation of charge within the bond itself.
The resulting separation of charge is represented by the Greek letter delta (\(\delta\)) to denote a partial charge. The atom with the stronger electron pull acquires a partial negative charge (\(\delta-\)), while the atom with the weaker pull acquires a partial positive charge (\(\delta+\)). A larger difference in the electronegativity values between the two atoms results in a more polar bond, meaning the partial charges are greater.
Hydrogen’s Standard Behavior in Chemical Bonds
Hydrogen is situated near the center of the Pauling electronegativity scale, with a value of approximately 2.20. This intermediate value explains why hydrogen’s charge depends so heavily on the specific element it is bonded to. In the vast majority of common chemical compounds, hydrogen bonds with elements that are substantially more electronegative than itself.
For instance, when hydrogen bonds with oxygen (3.44), nitrogen (3.04), fluorine (3.98), or chlorine (3.16), the shared electrons are pulled closer to the non-metal atom. This electron shift causes hydrogen to take on a partial positive charge (\(\delta+\)), a state that is fundamental to the chemistry of water (\(\text{H}_2\text{O}\)) and ammonia (\(\text{NH}_3\)). Even in methane (\(\text{CH}_4\)), where carbon’s electronegativity (2.55) is only slightly higher than hydrogen’s, hydrogen still carries the partial positive character.
Conditions for Hydrogen to Become Negative
The condition for hydrogen to gain a negative charge is straightforward: it must bond with an element that has a lower electronegativity value than its own 2.20. This reversal of polarity is observed in compounds known as hydrides, which are typically formed when hydrogen reacts with highly electropositive metals. The behavior of hydrogen in these compounds can result in either a partial or a full negative charge.
Partial Negative Charge (\(\delta-\))
A partial negative charge on hydrogen is found in certain covalent compounds where the bonding partner is less electronegative but the bond is still shared. This situation occurs with elements like Boron, which has a Pauling electronegativity of 2.04. In boron hydrides, such as diborane (\(\text{B}_2\text{H}_6\)), the electrons are pulled slightly toward the hydrogen atoms, giving them a partial negative charge (\(\delta-\)). In this context, the hydrogen acts as a source of electrons, which is the opposite of its behavior in water or ammonia.
Full Negative Charge (\(\text{H}^-\))
The most definitive example of negative hydrogen is the hydride ion, \(\text{H}^{-}\), which carries a full negative charge. This ion forms in ionic hydrides when hydrogen bonds with highly electropositive metals from Group 1 (alkali metals) or Group 2 (alkaline earth metals) of the periodic table. For example, in sodium hydride (\(\text{NaH}\)), sodium (electronegativity 0.93) essentially transfers its valence electron completely to the hydrogen atom. This complete electron transfer allows hydrogen to achieve the stable electron configuration of helium, resulting in a true ionic compound containing the negatively charged \(\text{H}^{-}\) ion.