Ionic and Covalent Bonding Study Pack

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Last updated May 22, 2026

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Ionic and Covalent Bonding Study Guide

Break down the key distinctions between ionic and covalent bonding, from electron transfer and electronegativity differences to shared electron pairs, bond order, and dipole moments. Covers how bond character shifts from nonpolar to polar to fully ionic, plus how ionic crystal lattice structures explain high melting points and conductivity — exactly what you need for bonding units.

Key Takeaways

  • •Ionic bonds form when one atom transfers electrons to another, creating oppositely charged ions that attract each other electrostatically, typically between metals and nonmetals with large electronegativity differences (greater than ~1.7).
  • •Covalent bonds form when two atoms share one or more pairs of electrons, typically between nonmetals with similar electronegativities, allowing each atom to approach a full valence shell.
  • •Electronegativity—an atom's tendency to pull shared electrons toward itself—determines bond character: small differences produce nonpolar covalent bonds, moderate differences produce polar covalent bonds, and large differences produce ionic bonds.
  • •Bond order (single, double, triple) reflects the number of shared electron pairs; higher bond order correlates with shorter bond length and greater bond energy.
  • •Polar covalent bonds create partial charges (δ+ and δ−) on bonded atoms, giving molecules a dipole moment that influences physical and chemical behavior.
  • •Ionic compounds form extended crystal lattice structures that result in high melting points, brittleness, and electrical conductivity when dissolved or melted.

Electronegativity: The Foundation of Bond Type

Before distinguishing ionic from covalent bonds, it helps to understand the underlying property that determines how electrons behave between two atoms: electronegativity.

Electronegativity Defined

  • •Electronegativity is a measure of how strongly an atom attracts electron density toward itself when it is part of a chemical bond.
  • •The Pauling scale is the most widely used system, assigning values from about 0.7 (francium) to 4.0 (fluorine).
  • •Electronegativity increases across a period (left to right) and decreases down a group on the periodic table, following the same trends as nuclear charge and atomic radius.

Electronegativity Difference and Bond Character

  • •When two identical or very similar atoms bond, the electronegativity difference (ΔEN) is near zero and electrons are shared equally — a nonpolar covalent bond.
  • •A ΔEN between roughly 0.4 and 1.7 results in unequal sharing of electrons — a polar covalent bond.
  • •A ΔEN greater than approximately 1.7 indicates that one atom dominates electron ownership strongly enough that chemists classify the bond as ionic, though this cutoff is a guideline, not an absolute rule.
  • •Bonds exist on a continuous spectrum from purely covalent to predominantly ionic; the categories are conceptually useful but not sharply divided.

Ionic Bonding: Electron Transfer and Ion Formation

Ionic bonding arises from the complete (or near-complete) transfer of one or more electrons from a metal atom to a nonmetal atom, producing oppositely charged ions that attract each other.

How Ions Form

  • •Metals, which have low ionization energies and low electronegativities, release valence electrons to form positively charged cations (e.g., Na loses one electron to become Na⁺).
  • •Nonmetals, which have high electron affinities and high electronegativities, accept those electrons to form negatively charged anions (e.g., Cl gains one electron to become Cl⁻).
  • •Both resulting ions achieve electron configurations matching the nearest noble gas, satisfying the octet rule.

Electrostatic Attraction and Lattice Structure

  • •The force holding ionic compounds together is electrostatic attraction between cations and anions — not a bond between one specific pair, but an omnidirectional attraction.
  • •Ionic compounds arrange into a crystal lattice, a three-dimensional repeating grid in which each ion is surrounded by multiple ions of opposite charge.
  • •The sodium chloride lattice, for example, places each Na⁺ surrounded by six Cl⁻ ions and vice versa.

Physical Properties Explained by Ionic Structure

  • •High melting and boiling points result from the large amount of energy needed to overcome lattice energy — the energy holding the ionic lattice together.
  • •Ionic solids are brittle because shifting the lattice even slightly aligns like charges next to each other, causing electrostatic repulsion that shatters the crystal.
  • •Ionic compounds conduct electricity only when dissolved in water or melted, because those processes free the ions to move and carry charge; solid ionic compounds do not conduct electricity because their ions are locked in fixed positions.

About this Study Pack

Created by Kibin to help students review key concepts, prepare for exams, and study more effectively. This Study Pack was checked for accuracy and curriculum alignment using authoritative educational sources. See sources below.

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