8.2 & 8.3: Electronegativity, Bond Polarity, and Dipole Moments

Electronegativity

• Definition: Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself. It describes the different affinities of atoms for the electrons within a chemical bond.

Linus Pauling's Method for Determining Electronegativity

• Developer: Linus Pauling (1901-1995), an American scientist and Nobel Prize laureate (Chemistry and Peace), developed the most widely accepted method for determining electronegativity values.

• Hypothetical Molecule Example: To understand Pauling's model, consider a hypothetical molecule $HX$.

• Relative Electronegativity Determination: The relative electronegativities of the $H$ and $X$ atoms are determined by comparing the measured $H-X$ bond energy with an "expected" $H-X$ bond energy.

• "Expected" Bond Energy Formula: The expected $H-X$ bond energy is calculated as the average of the $H-H$ and $X-X$ bond energies:
  $\text{Expected H-X bond energy} = \frac{\text{H-H bond energy} + \text{X-X bond energy}}{2}$

• Difference ($\Delta$) in Bond Energies: The difference, represented by $\Delta$, between the actual (measured) and expected bond energies is: $\Delta = (H-X)<em>{\text{act}} - (H-X)</em>{\text{exp}}$

  • Identical Electronegativities: If $H$ and $X$ have identical electronegativities, $(H-X)<em>{\text{act}}$ and $(H-X)</em>{\text{exp}}$ are the same, resulting in $\Delta = 0$. In this case, electrons are shared equally.

  • Different Electronegativities (Polarity): If $X$ has a greater electronegativity than $H$, the shared electron(s) will be pulled closer to the $X$ atom. This creates a polar molecule with a partial negative charge on $X$ ($\delta^-$) and a partial positive charge on $H$ ($\delta^+$):
    $H^{\delta^+} - X^{\delta^-}$

• Ionic Component and Bond Strength: This polar bond can be viewed as having both ionic and covalent components. The attraction between the partially (and oppositely) charged $H$ and $X$ atoms leads to a greater bond strength. Consequently, $(H-X)<em>{\text{act}}$ will be larger than $(H-X)</em>{\text{exp}}$.

• Magnitude of $\Delta$: The greater the difference in the electronegativities of the atoms, the larger the ionic component of the bond and the greater the value of $\Delta$. This $\Delta$ value allows for the assignment of relative electronegativities.

Periodic Trends in Electronegativity

• **General Trends (Figure 8.3):

  • Across a Period (Left to Right): Electronegativity generally increases.

  • Down a Group (Top to Bottom): Electronegativity generally decreases.

• Range of Values: Electronegativity values range from $\mathbf{4.0}$ for Fluorine ($F$) to $\mathbf{0.7}$ for Cesium ($Cs$).

• **Examples of Pauling Electronegativity Values (Excerpt from Figure 8.3):

  • $F: 4.0$

  • $O: 3.5$

  • $Cl: 3.0$

  • $N: 3.0$

  • $Br: 2.8$

  • $C: 2.5$

  • $S: 2.5$

  • $H: 2.1$

  • $Na: 0.9$

  • $Cs: 0.7$

Relationship Between Electronegativity and Bond Type (Table 8.1)

• Electronegativity Difference of Zero:

  • Bond Type: Covalent

  • Electron Sharing: Electrons are shared equally; no polarity develops.

  • Example: H-H bond.

• **Intermediate Electronegativity Difference (Increasing Polarity):

  • Bond Type: Polar covalent

  • Electron Sharing: Electrons are shared unequally.

  • Characteristics: Bonds have both covalent and ionic character.

• Large Electronegativity Difference:

  • Bond Type: Ionic

  • Electron Transfer: Electron transfer occurs, forming ions that constitute an ionic substance.

  • Characteristics: Predominantly ionic character.

Interactive Example 8.1: Relative Bond Polarities

• Task: Order the bonds $\text{H-H}$, $\text{O-H}$, $\text{Cl-H}$, $\text{S-H}$, and $\text{F-H}$ according to polarity.

• Principle: Bond polarity increases as the difference in electronegativity between the bonded atoms increases.

• **Electronegativity Values (from Figure 8.3):

  • $H: 2.1$

  • $S: 2.5$

  • $Cl: 3.0$

  • $O: 3.5$

  • $F: 4.0$

• **Calculated Electronegativity Differences and Polarity Order:

  1. $H-H$: ($2.1$ - $2.1$) = $0$

  2. $S-H$: ($2.5$ - $2.1$) = $0.4$

  3. $Cl-H$: ($3.0$ - $2.1$) = $0.9$

  4. $O-H$: ($3.5$ - $2.1$) = $1.4$

  5. $F-H$: ($4.0$ - $2.1$) = $1.9$

• **Polarity Order (Least to Most Polar): \text{H-H} < \text{S-H} < \text{Cl-H} < \text{O-H} < \text{F-H}

  • Electronegativity Difference: $0 \quad \quad 0.4 \quad \quad 0.9 \quad \quad 1.4 \quad \quad 1.9$

  • Bond Type: \text{Covalent bond} \quad \text{Polar covalent bond} \quad \text{Polarity increases}

Critical Thinking: Implications of Uniform Electronegativity

• Scenario: What if all atoms had the same electronegativity values?

• **Impact on Bonding:

  • Only Covalent Bonds: All bonds between atoms would be purely covalent (or nonpolar covalent).

  • No Ionic Bonds: No electron transfer would occur, preventing the formation of ionic substances.

  • No Polar Covalent Bonds: No partial charges ($\delta^+, \delta^-$) would develop, eliminating polar molecules.

• **Noticeable Differences:

  • Lack of Polarity-Dependent Properties: Many physical and chemical properties that depend on molecular polarity (e.g., solubility in polar solvents, boiling points due to dipole-dipole interactions, hydrogen bonding) would not exist or would be drastically altered.

  • Reduced Reactivity/Specificity: Chemical reactions relying on charge differential and electrostatic attractions would be significantly different or absent.

  • Uniformity in Electron Sharing: All shared electrons would be distributed equally between bonded atoms.