Water (\(\text{H}_2\text{O}\)) is a polar molecule, a fundamental property that dictates nearly all of its unique characteristics. Polarity means the molecule possesses an uneven distribution of electrical charge, creating a positive end and a negative end, much like a tiny bar magnet. This internal electrical asymmetry is called a dipole moment, and it arises directly from the specific chemical structure of the water molecule. Understanding this dipole moment requires looking closely at the bonds between the atoms and the resulting three-dimensional shape of the molecule.
Unequal Electron Sharing and Polar Bonds
The first step in water’s polarity involves how the oxygen and hydrogen atoms share their electrons within the covalent bonds. Atoms have a property called electronegativity, which is a measure of an atom’s ability to attract a shared pair of electrons toward itself. Oxygen exhibits a significantly higher electronegativity value (3.44) compared to hydrogen’s value (2.20).
This difference in electron-pulling power means that the oxygen atom pulls the electrons shared in the two O-H covalent bonds closer to itself. The shared electrons spend more time orbiting the oxygen atom than they do the hydrogen atoms. Because the electrons are pulled closer to the oxygen nucleus, the oxygen atom gains a slight negative electrical charge, denoted as partial negative (\(\delta-\)).
Conversely, the hydrogen atoms, having their electrons pulled away, become electron-deficient. This results in each hydrogen atom acquiring a slight positive electrical charge, denoted as partial positive (\(\delta+\)). These localized charge separations within each individual O-H bond are known as bond dipoles, establishing the necessary condition for a polar molecule.
The Crucial Role of Molecular Shape
While the partial charges in the O-H bonds are necessary, they are not sufficient to make the water molecule polar; the molecular shape is equally important. If the water molecule were linear, like carbon dioxide (\(\text{CO}_2\)), the two bond dipoles would pull in opposite directions, causing them to cancel each other out. This cancellation would result in a nonpolar molecule, despite having polar bonds.
However, the central oxygen atom in water possesses two pairs of electrons that are not involved in bonding, known as lone pairs. These lone pairs occupy space and exert a greater repulsive force on the two bonded hydrogen atoms. This electron-pair repulsion forces the two O-H bonds into a non-linear, or “bent,” geometry.
The resulting shape is bent, with a measured bond angle of approximately \(104.5^\circ\). Because the partial negative charge is concentrated on the oxygen end and the partial positive charges are on the hydrogen ends, this bent geometry prevents the bond dipoles from canceling. Instead, the two bond dipoles add together, resulting in a net electrical dipole moment for the molecule.
Practical Effects of Water’s Polarity
Water’s polarity has practical implications, most notably its ability to dissolve a vast number of substances, earning it the nickname “universal solvent.” The partial negative oxygen end of the water molecule is attracted to the positive ions or positive regions of other polar molecules. Simultaneously, the partial positive hydrogen ends are attracted to negative ions or negative regions of other molecules.
This charge-based attraction allows water to surround and pull apart the components of ionic compounds, such as salt, and other polar substances like sugars. The polarity also enables water molecules to engage in a powerful intermolecular force called hydrogen bonding. Hydrogen bonding is the attraction between the partial positive hydrogen of one water molecule and the partial negative oxygen of a neighboring water molecule.
This hydrogen bonding network gives water its high melting and boiling points, and contributes to its liquid properties. Cohesion, the tendency of water molecules to stick to one another, and adhesion, the tendency of water to stick to other polar surfaces, result from this strong intermolecular attraction. These properties are fundamental to life, allowing water to form drops, generate surface tension, and move through biological systems.