Electronegativity is a fundamental property of atoms that helps predict how elements behave when they combine to form molecules. Understanding the systematic changes in this property across a horizontal row, or period, is foundational to chemical science. These changes are governed by the underlying structure of the atom, which dictates the strength of the nucleus’s pull on its surrounding electrons. Examining this electron-attracting force makes it clear why the chemical characteristics of elements shift predictably across the periodic table.
What Electronegativity Measures
Electronegativity is a measure of an atom’s inherent ability to attract a shared pair of electrons toward itself when the atom is part of a chemical bond. This property allows chemists to predict the nature of a bond between two different atoms, distinguishing between purely covalent, polar covalent, or ionic interactions. The higher an atom’s electronegativity value, the more strongly it will pull the electron density of a shared bond closer to its own nucleus.
This atomic tendency is not a directly measurable quantity; instead, it is a calculated value derived from other properties. The most widely adopted system for quantifying this attraction is the Pauling scale, which assigns a numerical value to each element. On this scale, values range from a low of about 0.7 to a high of 4.0 for Fluorine. The difference between two bonded atoms’ electronegativity values determines the polarity of the resulting chemical bond.
The Observed Trend Across the Table
Electronegativity consistently increases as one moves from the left side to the right side across any given period of the periodic table. This means elements like the alkali metals (Group 1) exhibit the lowest values, while the halogens (Group 17) display the highest values. Elements on the far left, such as Sodium, have a weak attraction for a shared electron pair, while elements on the far right, such as Chlorine, have a very strong attraction.
This systematic increase across a period is a direct consequence of the changing atomic structure, which causes a shift from metallic character to non-metallic character. Elements on the left tend to have fewer valence electrons, making them more likely to lose electrons. Conversely, elements on the right already have a nearly full outer shell, making them more likely to attract electrons to complete their shell, which correlates with high electronegativity.
The Role of Effective Nuclear Charge
The primary physical force driving the increase in electronegativity across a period is the increasing effective nuclear charge, symbolized as \(Z_{eff}\). As one moves sequentially across a horizontal row, the number of protons in the nucleus increases steadily by one at each step. This addition of positive charge creates a stronger overall positive attraction for all surrounding electrons, including the valence electrons involved in bonding.
The effective nuclear charge represents the net positive charge experienced by the outermost valence electrons after accounting for the repelling effect of the inner-shell electrons. Since the number of protons is increasing without a significant change in the electron-blocking effect, the valence electrons are pulled inward more forcefully. This stronger inward pull results in a slight contraction of the atom’s size, meaning the atomic radius decreases across a period. A smaller atomic size places the valence electrons closer to the nucleus’s stronger positive charge, thereby magnifying the atom’s ability to attract a shared electron pair in a bond.
Why Shielding Stays Constant
The trend of increasing electronegativity is dominated by the nuclear charge because the electron shielding effect remains relatively consistent across a period. Electron shielding refers to the reduction of the nucleus’s attractive force on the outermost electrons due to the repulsion caused by the electrons in the occupied inner shells. When moving across a period, electrons are added to the same principal energy level, meaning they are all being placed into the same outermost shell.
Crucially, the number of inner, core electron shells that lie between the nucleus and the valence electrons does not change within a period. For instance, all elements in Period 3, from Sodium to Argon, have their valence electrons in the third energy level and are shielded only by the two filled inner shells. Because the count of these core, shielding electrons remains constant, the overall shielding effect they provide does not increase significantly. This allows the continuously increasing positive charge of the nucleus to exert a greater net pull on the valence electrons, reinforcing the observed trend of increasing electronegativity.