2.1 Polar Covalent Bonds and Electronegativity

Electronegativity determines how atoms share electrons in a covalent bond. When two atoms don't share equally, the bond becomes polar, and that polarity drives much of the reactivity you'll see throughout organic chemistry.

Polar Covalent Bonds and Electronegativity

Electronegativity and Polar Covalent Bonds

Electronegativity measures an atom's ability to attract shared electrons toward itself within a bond. Atoms like fluorine and oxygen have high electronegativity values and pull electrons strongly, while carbon and hydrogen have lower values and pull more weakly.

When two atoms with different electronegativities form a covalent bond, the electrons aren't shared equally. The more electronegative atom pulls electron density toward itself, developing a partial negative charge ($\delta-$), while the other atom is left with a partial positive charge ($\delta+$). This is a polar covalent bond.

The greater the electronegativity difference, the more polar the bond. Think of it as a spectrum:

• Nonpolar covalent (difference ≈ 0): Equal sharing. Example: C–C in ethane, where both atoms pull equally.
• Polar covalent (difference roughly 0.5–1.9): Unequal sharing. Example: O–H in water, where oxygen pulls electron density away from hydrogen.
• Ionic (difference ≥ 2.0): Electron transfer rather than sharing. Example: Na–Cl in sodium chloride.

The polarity of a bond can be quantified by its dipole moment, which measures the magnitude of charge separation multiplied by the distance between the charges.

Bond Polarity from Electronegativity Values

Electronegativity values come from the Pauling scale, the most widely used scale in organic chemistry. Two periodic trends to remember:

• Electronegativity increases from left to right across a period (e.g., Li → F).
• Electronegativity decreases going down a group (e.g., F → I).

To predict bond polarity, compare the electronegativity values of the two bonded atoms:

1. Look up each atom's Pauling electronegativity (common values: F = 4.0, O = 3.5, N = 3.0, C = 2.5, H = 2.1).
2. Calculate the difference.
3. Classify the bond:
   • Difference of 0 → nonpolar covalent
   • Difference between ~0.5 and 1.9 → polar covalent
   • Difference ≥ 2.0 → ionic

For example, in an O–H bond: $3.5 - 2.1 = 1.4$, which falls in the polar covalent range. Oxygen carries the $\delta-$ and hydrogen carries the $\delta+$. This uneven electron distribution creates regions of higher electron density around oxygen and lower electron density around hydrogen.

Electrostatic Potential Map Interpretation

Electrostatic potential maps (EPMs) are color-coded visuals of how electron density is distributed across a molecule's surface. The color scheme works like this:

• Red = electron-rich (partial negative charge). These regions have high electron density, often around lone pairs or electronegative atoms like oxygen.
• Blue = electron-poor (partial positive charge). These regions have low electron density, typically around less electronegative atoms like hydrogen bonded to O or N.
• Green/yellow = intermediate electron density.

The gradient from red → yellow → green → blue shows a smooth transition from most negative to most positive.

Why do EPMs matter for organic chemistry? They help you predict reactivity:

• Red regions are targets for electrophilic attack. Positively charged or electron-deficient species are attracted to these electron-rich areas.
• Blue regions are targets for nucleophilic attack. Negatively charged or electron-rich species are attracted to these electron-poor areas.

For instance, in a carbonyl group (C=O), the EPM shows red around oxygen and blue around carbon. That tells you nucleophiles will attack the carbon, which is exactly what happens in nucleophilic addition reactions.

Molecular Structure and Polarity

A molecule's overall polarity depends on both bond polarity and molecular geometry. Even if individual bonds are polar, the molecule as a whole can be nonpolar if the geometry causes the dipoles to cancel.

$CO_2$ is a classic example: each C=O bond is polar, but the molecule is linear (180°), so the two dipoles point in exactly opposite directions and cancel out. The result is a nonpolar molecule.

$H_2O$, by contrast, has a bent geometry (~104.5°). The two O–H bond dipoles point in roughly the same direction, so they reinforce each other. Water is a polar molecule with a net dipole moment.

To determine overall molecular polarity:

1. Identify all polar bonds and their directions ($\delta+$ → $\delta-$).
2. Determine the molecular geometry using VSEPR (valence shell electron pair repulsion).
3. Add the bond dipoles as vectors. If they cancel, the molecule is nonpolar. If they don't, it's polar.