Ionic and Covalent Bonding Study Pack
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Last updated May 27, 2026
Ionic and Covalent Bonding Study Guide
Break down the key distinctions between ionic and covalent bonding, from electron transfer and electronegativity thresholds to polar vs. nonpolar character, bond order, bond length, and the roles of lattice energy and bond dissociation energy in compound stability.
Key Takeaways
- •Ionic bonds form when one atom transfers one or more electrons to another, producing oppositely charged ions that attract each other electrostatically.
- •Covalent bonds form when two atoms share one or more pairs of electrons, allowing each atom to achieve a fuller valence shell without complete electron transfer.
- •Electronegativity difference between bonded atoms determines bond type: differences greater than ~1.7 generally indicate ionic character, while smaller differences indicate covalent character.
- •Covalent bonds are classified as nonpolar (equal sharing, electronegativity difference near 0) or polar (unequal sharing, partial charges on each atom), with polar covalent representing a continuum between purely nonpolar and fully ionic.
- •Bond order, bond length, and bond energy are interrelated: triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.
- •Lattice energy determines the stability of ionic compounds, while bond dissociation energy measures the strength of individual covalent bonds.
Electron Behavior as the Foundation of Chemical Bonding
All chemical bonds arise from the behavior of valence electrons — the electrons in an atom's outermost shell — and the drive of atoms toward more stable, lower-energy electron configurations.
Valence Electrons and the Octet Rule
- •Valence electrons occupy the highest principal energy level of an atom and are the only electrons involved in bonding.
- •The octet rule states that atoms tend to bond in ways that give them eight electrons in their valence shell, matching the stable configuration of noble gases.
- •Hydrogen is an exception, seeking only two valence electrons (a duet) to match the configuration of helium.
Electronegativity as the Deciding Factor
- •Electronegativity is a measure of how strongly an atom attracts shared electrons toward itself within a bond.
- •Linus Pauling developed the most widely used electronegativity scale, on which fluorine scores the highest value (3.98) and cesium scores among the lowest (0.79).
- •The difference in electronegativity between two bonded atoms predicts whether electrons will be transferred (ionic), unequally shared (polar covalent), or equally shared (nonpolar covalent).
Ionic Bonding: Electron Transfer and Electrostatic Attraction
Ionic bonding occurs when the electronegativity difference between two atoms is large enough that one atom strips electrons away from the other rather than sharing them, creating two oppositely charged particles that attract each other.
Formation of Cations and Anions
- •When a metal such as sodium (Na) loses a valence electron, it becomes a positively charged ion called a cation (e.g., Na⁺).
- •When a nonmetal such as chlorine (Cl) gains that electron, it becomes a negatively charged ion called an anion (e.g., Cl⁻).
- •The resulting electrostatic attraction between the cation and anion constitutes the ionic bond.
Ionic Compounds and Crystal Lattice Structure
- •Ionic compounds do not exist as discrete two-atom molecules; instead, each ion is surrounded by multiple ions of opposite charge in a repeating three-dimensional structure called a crystal lattice.
- •Lattice energy is the energy released when gaseous ions come together to form the ionic solid; higher lattice energy indicates a stronger, more stable ionic compound.
- •Lattice energy increases as ion charges increase and as ionic radii decrease, which is why MgO (Mg²⁺ and O²⁻) has a much higher lattice energy than NaCl (Na⁺ and Cl⁻).
Typical Properties of Ionic Compounds
- •Ionic compounds are typically hard, brittle crystalline solids with high melting points because overcoming the lattice energy requires substantial thermal energy.
- •They conduct electricity when dissolved in water or melted, because the ions become free to move and carry charge, but do not conduct in the solid state where ions are fixed in place.
Covalent Bonding: Electron Sharing Between Nonmetals
Covalent bonds form predominantly between nonmetal atoms that have similar or identical electronegativities, where pulling electrons away from either atom would require more energy than simply sharing them.
Single, Double, and Triple Bonds
- •A single bond consists of one shared pair of electrons (two electrons total) and is represented by one line between atoms, as in H–H or H–Cl.
- •A double bond consists of two shared pairs (four electrons) and a triple bond consists of three shared pairs (six electrons), as in O=O and N≡N respectively.
- •Bond order refers to the number of shared pairs: single bonds have order 1, double bonds order 2, and triple bonds order 3.
Bond Length and Bond Energy Trends
- •Bond length is the average distance between the nuclei of two bonded atoms; it decreases as bond order increases because additional shared electrons pull the nuclei closer together.
- •Bond dissociation energy is the energy required to break one mole of a specific bond homolytically in the gas phase; it increases with bond order.
- •For carbon–carbon bonds: C–C averages 347 kJ/mol and 154 pm; C=C averages 614 kJ/mol and 134 pm; C≡C averages 839 kJ/mol and 120 pm.
Lewis Structures as a Representation Tool
- •Lewis structures use dots to represent valence electrons and lines to represent shared pairs, allowing chemists to visualize how atoms satisfy the octet rule in a molecule.
- •Lone pairs are pairs of valence electrons on an atom that are not shared with another atom; they influence molecular shape even though they are not bonding electrons.
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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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What is the electronegativity value of fluorine on the Pauling scale?
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Electronegativity
Explain electronegativity in your own words. What does it measure, and how does the difference in electronegativity between two bonded atoms determine what type of bond forms?
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