Electronegativity Differences and Dipole Moments in Covalent Bonds
Electronegativity
Electronegativity is the ability of an atom to attract shared electrons in a covalent bond.
Think of it as a "tug-of-war" for electrons: the stronger the atom's pull, the higher its electronegativity.
Example- Fluorine (F), the most electronegative element, has a value of 4.0.
- Hydrogen (H), with a lower electronegativity of 2.2, exerts a weaker pull.
Electronegativity Differences and Bond Types
The difference in electronegativity ($\Delta \chi$) between two bonded atoms determines the type of bond they form:
- Nonpolar Covalent Bond ($\Delta \chi = 0$): Electrons are shared equally (e.g., $H_2$).
- Polar Covalent Bond ($0< \Delta \chi < 1.8$): Electrons are shared unequally, creating partial charges (e.g., $HCl$).
- Ionic Bond ($\Delta \chi >1.8$): Electrons are transferred completely, forming ions (e.g., $NaCl$).
Classifying Bond Types
- Carbon–Hydrogen (C–H):
- Electronegativity of C = 2.6, H = 2.2.
- $\Delta \chi = |2.6 - 2.2| = 0.4$, so the bond is nonpolar covalent.
- Hydrogen–Fluorine (H–F):
- Electronegativity of H = 2.2, F = 4.0.
- $\Delta \chi = |4.0 - 2.2| = 1.8$, so the bond is highly polar covalent.
Electronegativity differences provide a guideline, but remember that bonding exists on a continuum rather than in strict categories.
HintElectronegativity values can be found in Section 9 of the Data Booklet.
Bond Polarity and Partial Charges (δ+ and δ−)
Bond polarity
Bond polarity refers to the unequal sharing of electrons between two atoms in a covalent bond, resulting in a partial positive charge on one atom and a partial negative charge on the other.
- When two atoms in a bond have different electronegativities, the shared electrons spend more time near the more electronegative atom.
- This creates:
- A partial negative charge ($ \delta^- $) on the more electronegative atom.
- A partial positive charge ($ \delta^+ $) on the less electronegative atom.
In $HCl$:
- Chlorine ($ \chi = 3.2 $) attracts the electrons more strongly, becoming $ \delta^- $.
- Hydrogen ($ \chi = 2.2 $) becomes $ \delta^+ $.
Representing Bond Polarity
Bond polarity can be visualized in two ways:
- Partial Charges ($ \delta^+ $ and $ \delta^- $):
