U2 Chem- 2.26

Ionic and Covalent Bonds

Ionic Bonds

• Example: Lithium chloride (LiCl) is a common ionic compound that illustrates the nature of ionic bonding.

• In ionic bonds, electrons are transferred between atoms rather than shared. This typically occurs between metals and nonmetals.

• In the case of lithium (Li), it has one electron in its outer shell (valence shell), which it readily loses, resulting in a positively charged lithium ion (Li⁺).

• Conversely, chlorine (Cl) has seven electrons in its outer shell and needs one more to complete its octet. It gains the electron lost by lithium, forming a negatively charged chloride ion (Cl⁻).

• The electrostatic attraction between these oppositely charged ions creates a strong ionic bond, forming a stable ionic compound.

Electronegativity and Polar Covalent Bonds

• Electronegativity Values: Electronegativity is a measure of an atom's ability to attract and hold onto electrons. Here are the electronegativity values for two relevant elements:

  • Nitrogen: 3.0

  • Hydrogen: 2.1

• Delta Electronegativity (ΔEN): This value helps to determine the type of bond between two atoms. For example, for the N-H bond:

  • ΔEN for N-H = 3.0 - 2.1 = 0.9

  • Since the ΔEN is between 0.5 and 1.9, this indicates a polar covalent bond, where electrons are shared unequally.

  • In this bond, nitrogen, being more electronegative, attracts electrons from hydrogen more frequently, resulting in a partial positive charge (δ+) on hydrogen and a partial negative charge (δ-) on nitrogen.

Cutoffs for Bond Type

• The type of bond formed between two atoms can be classified based on the difference in their electronegativities:

  • 0.5 to 1.9: Polar covalent bond, where there is unequal sharing of electrons.

  • Above 1.9: Ionic bond, where electrons are transferred from one atom to another.

  • Below 0.5: Nonpolar covalent bond, where electrons are shared equally between atoms, resulting in no charge separation.

Visualization of Charge Distribution

• Covalent Bonding: A classic example is the hydrogen molecule (H₂), which has an electronegativity difference (ΔEN) of 0, resulting in an even distribution of electrons. This molecule is considered to have a pure covalent bond due to the equal sharing of electrons.

• Example of HF (Hydrogen Fluoride):

  • Electronegativity Values: Fluorine has a high electronegativity of 4.0, while hydrogen's is 2.1.

  • The ΔEN for HF = 4.0 - 2.1 = 1.9, indicating a highly polar bond that approaches ionic character due to the significant difference in electronegativity.

  • Charge distribution shows fluorine acquiring a partial negative charge (δ-) and hydrogen a partial positive charge (δ+), highlighted in diagrams where fluorine appears in red to indicate its negative charge, and hydrogen is relatively positive in color-coded models.

Understanding these concepts is critical for comprehending chemical reactions, molecular structure, and the properties of substances in chemistry.