Electronegativity is a fundamental property of atoms that governs how they interact to form chemical bonds. It describes an atom’s inherent ability to attract shared electrons toward itself when chemically bonded to another atom. This concept is central to understanding the nature of chemical bonds, which determines the structure and behavior of all molecules and compounds. The variation in this electron-attracting power among elements explains the diverse chemical and physical properties observed in substances.
Defining the Electron Tug-of-War
Electronegativity measures an atom’s pulling power on the electron cloud it shares with another atom in a bond. This attraction is only relevant when atoms are linked together, making it a property of an atom within a molecule, not an isolated atom. The scenario is often compared to an atomic tug-of-war, where the rope represents the shared pair of valence electrons. If the two bonded atoms were identical, the electrons would be shared equally, resulting in a balanced pull.
When atoms of different elements bond, the pull is unequal, causing the shared electron density to shift closer to the nucleus of the stronger atom. This pulling strength is influenced by the atom’s nuclear charge and its atomic radius. A greater number of protons increases the positive charge, creating a stronger attractive force on the shared electrons. Conversely, a larger atomic size means valence electrons are farther from the nucleus, which weakens the nuclear attraction and lowers the atom’s electron-pulling power.
The Electronegativity Scale
To quantify this electron-attracting ability, American chemist Linus Pauling developed the most widely used scale in 1932. The Pauling scale is a dimensionless system, meaning its values do not have units, and it is based on relative bond energies. The difference in electronegativity between two atoms is proportional to the extra stability of the bond they form.
The values on this scale range from approximately 0.7 to 4.0. Fluorine, the element with the strongest ability to attract electrons, is assigned the highest value of 4.0. Elements like Cesium and Francium, which readily give up electrons, have the lowest values, around 0.7 to 0.8. These numbers are used to compare the electron-attracting power of different elements.
Patterns on the Periodic Table
Electronegativity values follow predictable patterns on the periodic table, explained by the underlying atomic structure. Moving from left to right across any period, electronegativity generally increases. This occurs because the number of protons in the nucleus increases while valence electrons are added to the same energy level. The resulting increase in nuclear charge creates a stronger net pull on the shared electrons.
Conversely, moving down a group, the electronegativity generally decreases. As the atomic number increases, electrons fill new, larger energy levels. This causes the atomic radius to increase, placing the outermost bonding electrons farther away from the nucleus. This increased distance, combined with the shielding effect of inner electrons, weakens the nucleus’s ability to attract external electrons.
The Role in Determining Chemical Bonds
Electronegativity predicts the type of chemical bond that will form between two atoms. The absolute difference in the electronegativity values (\(\Delta EN\)) of the two bonded atoms determines the character of the bond. Bond types exist along a continuous spectrum, ranging from equal sharing to complete electron transfer, but they are typically categorized into three main groups.
When the electronegativity difference is small (usually less than 0.4), the electrons are shared almost equally, resulting in a nonpolar covalent bond. Examples include the bond between two oxygen atoms in an \(\text{O}_2\) molecule (\(\Delta EN\) is zero) or the bond in methane (\(\text{C-H}\)). If the difference is moderate (typically between 0.4 and 1.7), a polar covalent bond forms. In water (\(\text{H}_2\text{O}\)), the higher electronegativity of oxygen pulls the shared electrons closer, creating a slight negative charge (\(\delta-\)) on oxygen and a slight positive charge (\(\delta+\)) on the hydrogen atoms.
A large difference in electronegativity (generally greater than 1.7) causes the more electronegative atom to strip the electron completely from the less electronegative atom. This complete transfer results in the formation of positively and negatively charged ions, held together by the electrostatic attraction of an ionic bond. Sodium chloride (\(\text{NaCl}\)) is a common example, where the large \(\Delta EN\) leads to the transfer of an electron and the formation of \(\text{Na}^+\) and \(\text{Cl}^-\) ions.