Chlorine is more electronegative than bromine. This distinction is a fundamental concept in chemistry, as electronegativity governs how atoms interact and form chemical bonds. The relative ability of chlorine and bromine to attract electrons influences the nature of the compounds they form, affecting molecular polarity and chemical reactivity. Understanding this difference requires examining what electronegativity represents and how atomic structure influences this characteristic.
What Electronegativity Represents
Electronegativity is a measure of an atom’s tendency to attract a shared pair of electrons toward itself when participating in a chemical bond. This property is a relative scale indicating the pull an atomic nucleus exerts on bonding electrons. Atoms with a high electronegativity value exert a stronger pull on shared electrons.
Linus Pauling developed the most commonly used scale to quantify this property, assigning values that typically range from 0.7 to 4.0. The difference in electronegativity between two bonded atoms determines the polarity of the bond. A large difference means the electrons are shared unequally, creating a polar bond where electron density shifts toward the more electronegative atom.
How Atomic Structure Determines the Difference
The difference in electronegativity between chlorine and bromine stems from their positions on the periodic table and their atomic structures. Both are halogens in Group 17, but chlorine sits above bromine. Electronegativity generally decreases as you move down a group.
This decrease occurs because of the increase in atomic size, specifically the number of electron shells. Chlorine atoms have three electron shells, while bromine atoms have four. The additional electron shell in bromine means its valence electrons are farther from the positively charged nucleus.
The greater distance weakens the nucleus’s effective attraction on shared electrons. Furthermore, bromine’s extra inner electron shells create a stronger “shielding effect.” These inner electrons partially block the nucleus’s attractive pull.
Chlorine, having fewer inner shells and a smaller atomic radius, experiences less shielding. This allows its nucleus to exert a stronger attractive force on shared electrons. This structural difference results in measurable values on the Pauling scale, confirming chlorine’s stronger pull.
Chlorine’s electronegativity is approximately 3.16, while bromine’s is lower at 2.96. This difference means that in a bond between the two elements, shared electrons are more attracted to the chlorine atom. The atomic structure—fewer electron shells and less shielding—is the mechanism making chlorine the more electronegative element.