Electronegativity is a fundamental chemical property describing an atom’s inherent ability to attract electrons toward itself within a chemical bond. When two different atoms bond, their electron-attracting powers are rarely equal, leading to an unequal distribution of shared electrons. This disparity is quantified by calculating the electronegativity difference (\(\Delta\)EN). The resulting difference value is a powerful predictive tool, allowing chemists to determine the nature and polarity of the chemical bond.
What Electronegativity Measures
Electronegativity is a relative scale, not an absolute, measurable energy like ionization energy. The most widely adopted system is the Pauling scale, developed by Linus Pauling. Pauling established the numerical scale by setting the value for fluorine, the element with the highest electron affinity, at 4.0. All other elements were assigned values relative to this maximum, resulting in a scale where all elements fall between 0.7 and 4.0.
These assigned values follow predictable patterns across the periodic table. Electronegativity increases as you move from left to right across any period because the rising number of protons creates a stronger pull on the valence electrons. Conversely, the values decrease as you move down a group. Moving down a group adds more electron shells, which increases the atomic radius and shields the outer electrons, weakening the attractive force. The highest electronegativity elements are found in the upper right section, while the lowest values belong to elements in the lower left, such as Francium (0.7).
Step-by-Step Calculation
The process for finding the electronegativity difference is a straightforward exercise in subtraction. The first step involves identifying the two elements forming the chemical bond. Next, locate the Pauling electronegativity value for each atom using a reliable chart or periodic table. For example, in hydrogen chloride (HCl), the value for Hydrogen (H) is 2.1 and the value for Chlorine (Cl) is 3.0.
The final step is to calculate the absolute difference between these two values. This is achieved by subtracting the smaller value from the larger one, ensuring the result is always positive. Using the HCl example: Difference = 3.0 (Cl) – 2.1 (H) = 0.9. This number represents the magnitude of the difference in electron-attracting power between the two atoms.
Classifying Bond Types Based on the Difference
The practical value of calculating the electronegativity difference lies in its ability to predict the type of bond that forms between the two atoms. Chemists use the calculated number as a guideline to classify the bond into one of three main categories, reflecting a continuum of electron sharing.
Nonpolar Covalent Bonds
When the electronegativity difference is small (0 to 0.4), the bond is classified as Nonpolar Covalent. In this bond type, electrons are shared almost equally between the two atoms, such as in a bond between two identical atoms like O\(_{2}\), where the difference is zero.
Polar Covalent Bonds
As the difference increases (0.4 to 1.7), the bond becomes Polar Covalent. This intermediate range signifies unequal sharing, where the atom with the higher electronegativity value pulls the electron density closer to itself. This unequal pull creates a partial negative charge on the more attractive atom and a partial positive charge on the less attractive atom.
Ionic Bonds
A large difference (greater than 1.7) indicates an Ionic bond. A difference this large means the more electronegative atom has such a strong pull that it strips the electron completely away from the less electronegative atom, leading to the formation of positive and negative ions. These numerical cutoffs are general guides, as bonding is a smooth transition from equal sharing to complete electron transfer.