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.
- 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.
Electronegativity 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^- $):
- Example: $ H^{\delta +} - Cl^{\delta -} $.
- Dipole Moment Vector ($\vec{μ}$):
- The arrow starts with a “+” sign at the positive end and points toward the negative end.
- Example: $ H \rightarrow Cl $.
Dipole moment vector
A dipole moment vector is a vector pointing from the less electronegative atom to the more electronegative atom.
When drawing dipole moment vectors, always use molecular geometry to determine the direction and magnitude of the dipoles.
Electronegativity and Bond Polarity
Dipole Moments: Measuring Bond Polarity
Dipole moment
A dipole moment describes the separation of electrical charge in a bond or molecule due to differences in electronegativity.
It results in partial positive and negative regions, making the molecule polar.
In water $(H_2O)$, oxygen attracts electrons more strongly than hydrogen, creating a dipole.
A larger dipole moment indicates a greater charge separation, influencing properties like solubility and boiling point.
Molecular Dipole Moment vs. Bond Dipoles
While individual bonds may have dipoles, the overall molecular dipole depends on the geometry of the molecule:
- Polar Molecules: Bond dipoles do not cancel out, resulting in a net dipole moment (e.g., $ H_2O $).
- Nonpolar Molecules: Bond dipoles cancel out due to symmetry, resulting in no net dipole moment (e.g., $ CO_2 $).
Water vs. Carbon Dioxide
- Water ($ H_2O $):
- Geometry: Bent.
- The bond dipoles of the two $ O–H $ bonds do not cancel, resulting in a net dipole moment.
- -> Polar molecule.
- Carbon Dioxide ($ CO_2 $):
- Geometry: Linear.
- The bond dipoles of the two $ C=O $ bonds cancel each other out.
- -> Nonpolar molecule.
Use VSEPR theory to predict molecular geometry and determine whether bond dipoles will cancel.
- Assuming all molecules with polar bonds are polar:
- Always consider molecular geometry.
- Symmetrical molecules (e.g., $ CO_2 $) can have polar bonds but still be nonpolar overall.
- Forgetting to calculate electronegativity differences:
- Always use $ \Delta \chi $ to determine bond polarity and compare it to the bonding continuum.
Bond DIpole and Molecular Dipole
- Calculate the electronegativity difference for the bonds in $ HF $, $ CH_4 $, and $ NH_3 $. Classify each bond as nonpolar, polar, or ionic.
- Draw the dipole moment vector for $ HCl $ and $ CO_2 $.
- Explain why $ CHCl_3 $ is polar but $ CCl_4 $ is nonpolar.
