10.3 Electronegativity and Polarity

Definition and Core Concepts of Electronegativity

• Definition: Electronegativity is a measure of how strongly an atom attracts bonding electrons to itself when involved in a chemical bond.

• Atomic Perspective: Smaller atoms tend to have higher electronegativity values. This occurs because their valence electrons are closer to the nucleus, allowing the nucleus to exert a stronger attractive force on those outer electrons.

• Metallic vs. Nonmetallic Trends:

  • Nonmetals: Generally have high electronegativities because they have a high affinity for gaining electrons to achieve stability.

  • Metals: Generally have low electronegativities because they naturally tend to lose electrons to achieve stability.

• Element Extremes: Fluorine is identified as the most electronegative element on the periodic table.

Periodic Trends for Electronegativity

• Horizontal Trend (Across a Period): In general, electronegativity increases when moving from left to right across each period. This is because atoms are getting smaller and the nucleus becomes more effective at attracting electrons due to the decrease in distance between the nucleus and the valence shell.

• Vertical Trend (Up a Group): Electronegativity increases going up a group. Similar to the horizontal trend, atoms get smaller as you move up a group, resulting in a stronger attraction between the nucleus and outer electrons.

• Sublevel Focus: For fundamental chemistry studies, focus is primarily placed on elements containing valence electrons in the $s$-sublevel and the $p$-sublevel.

• Transition Metals: While electronegativity values exist for transition metals, they are harder to predict and are generally considered less critical for introductory molecular structure analysis.

Bond Polarity and Classification

• Bonding Mechanisms:

  • Ionic Bonding: Involves the complete transfer of electrons from one atom to another.

  • Covalent Bonding: Involves the sharing of electrons between atoms. This sharing can be equal or unequal.

• Unequal Sharing ("Hogging"): In many covalent bonds, one atom attracts the electrons more aggressively than the other. This results in a "polar" bond.

• Determining Bond Type by Electronegativity Difference ($\Delta EN$): To classify a bond, subtract the electronegativity of the less electronegative element from the more electronegative element (always subtract the smaller value from the larger value to get a positive difference).

Examples of Bond Polarity Calculations

• Example #1: The bond in $H_2$

  • Step 1 (Values): $H = 2.1$

  • Step 2 (Subtraction): $2.1 - 2.1 = 0$

  • Step 3 (Classification): The difference is $0$, meaning the bond is non-polar covalent. Electrons are shared exactly equally.

• Example #2: The bond in $HCl$

  • Step 1 (Values): $H = 2.1$, $Cl = 2.9$

  • Step 2 (Subtraction): $2.9 - 2.1 = 0.8$

  • Step 3 (Classification): The difference is $0.8$, which falls into the polar covalent range. The $Cl$ atom "hogs" the electrons because it possesses the higher electronegativity.

• Comparative Polarity:

  • $H - Br$: $2.8 - 2.1 = 0.7$

  • $H - Cl$: $2.9 - 2.1 = 0.8$

  • Comparing the two, the $HCl$ bond is more polar than the $HBr$ bond because it has a larger difference in electronegativity.

Molecular Polarity and Symmetry

• Dipole Definition: A dipole is the special name given to the uneven sharing of electrons within a bond or molecule.

• Determining Molecular Polarity: This is determined by two factors: bond polarity and molecular shape (geometry).

  • Rule 1: If all bonds in a molecule are non-polar, the entire molecule is non-polar.

  • Rule 2: If the bonds are polar, the symmetry of the molecule determines the final polarity.

• Cancellation of Dipoles: If polar bonds pull in equal and opposite directions (symmetrical arrangement), they cancel each other out, resulting in a non-polar molecule.

• General Heuristic:

  • Asymmetrical = Polar

  • Symmetrical = Non-polar

Comparison Case Study: $H_2O$ vs. $BeH_2$

• Water ($HOH$ or $H_2O$):

  • Bond Polarity: $H - O$ is $3.5 - 2.1 = 1.4$ (Polar bond).

  • Molecular Geometry: The presence of two sets of unpaired (lone) electrons on the oxygen atom forces a tetrahedral electron arrangement, resulting in a bent molecular shape.

  • Result: Because it is bent (asymmetrical), the dipoles do not cancel. The molecule is polar.

• Beryllium Hydride ($HBeH$):

  • Bond Polarity: $H - Be$ is $2.1 - 1.5 = 0.6$ (Polar bond).

  • Molecular Geometry: There are no lone pairs on the central Beryllium atom. The molecule assumes a linear shape.

  • Result: Because it is linear and symmetrical, the dipoles pull in opposite directions and cancel. The molecule is non-polar.

Practice Problems

1. Which of the following molecules contain polar bonds?

• a. $ClF_5$: $F(4.1) - Cl(2.9) = 1.2$ (Polar bond)

• b. $ClO_2^-$: $O(3.5) - Cl(2.9) = 0.6$ (Polar bond)

• c. $TeCl_4^{2-}$: $Cl(2.9) - Te(2.0) = 0.9$ (Polar bond)

• d. $PCl_3$: $Cl(2.9) - P(2.1) = 0.8$ (Polar bond)

• e. $O_2$: $O(3.5) - O(3.5) = 0$ (Non-polar bond)

• f. $PH_2^+$: $P(2.1) - H(2.1) = 0$ (Non-polar bond)

2. Determine if the following molecules are polar or non-polar:

• a. $CS_2$:

  • Bonds: $C(2.5) - S(2.4) = 0.1$ (Non-polar bonds).

  • Molecular Result: Non-polar.

• b. $PCl_3$:

  • Bonds: $Cl(2.9) - P(2.1) = 0.8$ (Polar bonds).

  • Shape: Trigonal Pyramidal (Asymmetrical).

  • Molecular Result: Polar.

• c. $H_2O$:

  • Bonds: $O(3.5) - H(2.1) = 1.4$ (Polar bonds).

  • Shape: Bent (Asymmetrical).

  • Molecular Result: Polar.

• d. $NH_4^+$:

  • Bonds: $N(3.1) - H(2.1) = 1.0$ (Polar bonds).

  • Shape: Tetrahedral (Symmetrical).

  • Molecular Result: Non-polar.