Ionic and Covalent Bonding, Electronegativity, Molecular Polarity, and Intermolecular Forces

Ionic and Covalent Bonding

Formation of Ions and Ionic Compounds

Atoms achieve stable electron configurations by gaining or losing electrons to form ions. Metals tend to lose electrons and form positively charged ions (cations), while nonmetals tend to gain electrons and form negatively charged ions (anions). Ionic compounds are formed when cations and anions attract each other through electrostatic forces, resulting in a neutral compound.

• Cations: Formed by metals losing electrons (e.g., Na+, Ca2+).

• Anions: Formed by nonmetals gaining electrons (e.g., Cl-, O2-).

• Example: Sodium (Na) loses one electron to form Na+; chlorine (Cl) gains one electron to form Cl-. Together, they form NaCl.

Key Properties of Ionic Compounds:

• High melting and boiling points

• Conduct electricity when molten or dissolved in water

• Form rigid, crystalline structures

Covalent Bonding and Molecular Compounds

Covalent compounds are formed when nonmetal atoms share pairs of electrons to achieve stable electron configurations. These shared electrons create covalent bonds, resulting in discrete molecules.

• Single, double, or triple bonds may form depending on the number of shared electron pairs.

• Example: In H2O, each hydrogen shares one electron with oxygen, forming two single bonds.

• Properties: Lower melting/boiling points, poor electrical conductivity, often soft or brittle.

Compounds with Both Ionic and Covalent Bonds

Some compounds contain both ionic and covalent bonds, typically when a metal cation is bonded to a polyatomic ion (e.g., NaNO3, CaCO3), or when a polyatomic cation is present (e.g., NH4Cl). The ionic bond exists between the ions, while the atoms within the polyatomic ion are held together by covalent bonds.

Electronegativity and Bond Polarity

Electronegativity Trends

Electronegativity (EN) is the ability of an atom to attract shared electrons in a chemical bond. EN increases from left to right across a period and from bottom to top within a group. Fluorine is the most electronegative element.

• Metals: Low EN (tend to lose electrons)

• Nonmetals: High EN (tend to gain electrons)

Bond Polarity

The difference in electronegativity (ΔEN) between two bonded atoms determines the bond type:

• ΔEN = 0–0.4: Nonpolar covalent (electrons shared equally)

• ΔEN = 0.5–1.8: Polar covalent (electrons shared unequally; partial charges develop)

• ΔEN > 1.8: Ionic (electrons transferred; ions form)

Example: In H2O, O–H bonds are polar covalent (ΔEN = 1.4), with oxygen carrying a partial negative charge (δ−) and hydrogen a partial positive charge (δ+).

Molecular Polarity and VSEPR Theory

Determining Molecular Polarity

A molecule is polar if it contains polar bonds and the molecular shape does not allow the bond dipoles to cancel. The VSEPR theory (Valence Shell Electron Pair Repulsion) is used to predict molecular shapes based on regions of high electron density (bonds and lone pairs) around the central atom.

• Nonpolar molecules: All bond dipoles cancel due to symmetry (e.g., CO2, CCl4).

• Polar molecules: Bond dipoles do not cancel (e.g., H2O, NH3).

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are attractions between molecules or ions that influence physical properties such as melting and boiling points.

• Ionic bonds: Strongest; between ions in ionic solids (e.g., NaCl).

• Hydrogen bonds: Strong dipole–dipole attractions between H bonded to N, O, or F and a lone pair on another N, O, or F atom (e.g., H2O, NH3).

• Dipole–dipole attractions: Between polar molecules (e.g., HCl, SO2).

• Dispersion forces: Weakest; present in all molecules, but only force in nonpolar molecules (e.g., CH4, I2).

Hydrogen Bonding in Water

Hydrogen bonds are responsible for many unique properties of water, such as its high boiling point and the fact that ice floats. In water, the δ+ hydrogen of one molecule is attracted to the δ− oxygen of another molecule.

Applications to Health Sciences

• "Like dissolves like": Polar and ionic substances dissolve in water; nonpolar substances (fats, oils) do not. This principle explains why vitamins and drugs have different solubilities in body fluids.

• Evaporation of sweat: When sweat evaporates, hydrogen bonds between water molecules must be broken, absorbing heat and cooling the body.

• Transport of fats: Fats are nonpolar and do not dissolve in blood plasma (mostly water). The body uses lipoproteins (with polar surfaces and nonpolar cores) to transport fats.

Summary Table: Intermolecular Forces

Force	Particles	Strength	Examples
Ionic bonds	Ions	Strongest	NaCl, KBr
Hydrogen bonds	Polar molecules with N–H, O–H, F–H	Strong	H2O, NH3, ethanol
Dipole–dipole	Polar molecules	Moderate	HCl, SO2
Dispersion	Nonpolar molecules	Weakest	CH4, CO2, Cl2

Key Equations and Concepts

• Electronegativity difference: $\Delta EN = |EN_1 - EN_2|$

• Bond type classification:

  • $0 \leq \Delta EN \leq 0.4$: Nonpolar covalent

  • $0.5 \leq \Delta EN \leq 1.8$: Polar covalent

  • $\Delta EN > 1.8$: Ionic