Types of Chemical Bonds
Compare sharing, ionic character and delocalized metallic electrons using properties as evidence.
AP Chemistry Topic 2.1 · CED printed p. 44, 45 / PDF p. 51, 52. This teaching, practice and feedback were added for the AP edition; original author material remains separately traceable.
Before you start
- Electronegativity trends
- Valence electrons and charge
Sharing can be equal or unequal
A bond is stabilized by electrostatic attractions involving nuclei and valence electrons. Electronegativity describes how strongly a bonded atom attracts shared electron density. Across representative-element periods, rising effective nuclear attraction generally increases electronegativity; down groups, larger valence distance and shielding generally decrease it. This is an explanatory model, not a list of bond energies.
Two atoms with the same or similar electronegativities share relatively evenly, giving a nonpolar covalent bond. C–H is treated as effectively nonpolar for this AP model despite a small difference. In a polar covalent bond, the more electronegative atom has partial negative charge, δ−, and the other has δ+. These partial charges are not whole-ion charges. For comparable single bonds, a larger electronegativity difference generally gives a larger bond dipole; bond length also enters the dipole moment, so an electronegativity difference alone is not a complete numerical dipole calculation.
Bonding is a continuum, and properties are evidence
A strong shift of electron density toward one partner gives substantial ionic character. The ionic–covalent distinction is a useful model on a continuum, not a universal cutoff at one exact electronegativity difference. Metal–nonmetal combinations are often ionic, and nonmetal–nonmetal combinations are often covalent. A formula and measured properties together provide a stronger diagnosis: a brittle solid that conducts when molten but not when solid supports mobile-ion behavior on melting.
In a metal, valence electrons are delocalized across the solid rather than assigned to one atom or a two-atom bond. Mobile electrons explain solid-state conductivity, while cohesion remains as metal cores rearrange. Do not rank ionic, polar covalent and nonpolar covalent bond energies merely by electronegativity difference. Polarity describes distribution of charge; bond dissociation energy describes the energy required to separate a specified bonded pair.
Let's work through it
Work it through
Compare H–H, H–Cl and NaCl using supplied electronegativities H = 2.2, Cl = 3.2, Na = 0.9; then use conductivity data to check the NaCl model.
- H–H has ΔEN = 0 and equal sharing: nonpolar covalent. H–Cl has ΔEN = 1.0; Cl attracts the shared density more strongly, so Hδ+–Clδ− is polar covalent.
- NaCl combines a metal and nonmetal and is well modeled as Na+/Cl− in an extended ionic structure. The large ΔEN = 2.3 supports ionic character, but the number is not the only evidence.
- If NaCl does not conduct as a solid but conducts when molten, the phase-dependent availability of mobile ions supports the ionic model.
- No bond-energy ranking follows solely from 0, 1.0 and 2.3. Such a ranking needs specified bond-energy data or an appropriate interaction comparison.
Try it from memory
Pause & retrieve
Try it before you peek.
Write your answer and the reason from memory. Then compare with the worked answer. Your self-assessment is a study decision, not an automatically graded score.
Question 1
Mark partial charges for an O–H bond.
See every recall question with its answer
These references preserve each prompt and its answer for direct evidence links. Try the interactive questions above before opening the worked reasoning.
Retrieval 1
Mark partial charges for an O–H bond.
Worked answer and reason
Oδ−–Hδ+ because O attracts shared electron density more strongly.
The charges describe unequal sharing, not full electron transfer in an isolated O–H bond.
Retrieval 2
What would solid conductivity suggest about a sample believed to be a metal?
Worked answer and reason
It is consistent with delocalized mobile electrons.
A metallic model provides a charge carrier in the solid state.
Catch the tempting mistake
- Any ΔEN above a memorized cutoff proves ionic bonding.
- Bonding is a continuum; composition and observed properties support the most useful model.
- A more polar bond is always stronger.
- Dipole magnitude and dissociation energy are different properties.
Put it to work
Give the chemical reason as well as the calculation. Worked solutions are on a separate page.
- Unit 2 problem 1 · ap application · multiple
- Unit 2 problem 2 · ap synthesis · macroscopic
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For the teacher
- Ask students to identify charge carriers for each conductivity claim.
- Use both a trend explanation and a property-based diagnosis.
Source objectives and scope
Source edition: ap-chemistry-ced-fall2024-local-b5dfe867. See this topic in the course outline →
- 2.1.A
- Explain the relationship between the type of bonding and the properties of the elements participating in the bond.
- 2.1.A.1
- Electronegativity values for the representative elements increase going from left to right across a period and decrease going down a group. These trends can be understood qualitatively through the electronic structure of the atoms, the shell model, and Coulomb’s law.
- 2.1.A.2
- Valence electrons shared between atoms of similar electronegativity constitute a nonpolar covalent bond. For example, bonds between carbon and hydrogen are effectively nonpolar even though carbon is slightly more electronegative than hydrogen.
- 2.1.A.3
- Valence electrons shared between atoms of unequal electronegativity constitute a polar covalent bond. i. The atom with a higher electronegativity will develop a partial negative charge relative to the other atom in the bond. ii. In single bonds, greater differences in electronegativity lead to greater bond dipoles. iii. All polar bonds have some ionic character, and the difference between ionic and covalent bonding is not distinct but rather a continuum.
- 2.1.A.4
- The difference in electronegativity is not the only factor in determining if a bond should be designated as ionic or covalent. Generally, bonds between a metal and nonmetal are ionic, and bonds between two nonmetals are covalent. Examination of the properties of a compound is the best way to characterize the type of bonding.
- 2.1.A.5
- In a metallic solid, the valence electrons from the metal atoms are considered to be delocalized and not associated with any individual atom.
Science references for this teaching
- AP Chemistry Course and Exam Description, effective Fall 2024 (supplied source) — Exact topic-numbered learning objectives, essential knowledge, science practices and AP exclusions; supplied PDF SHA256 b5dfe8677ef3d88c613865d2e2a3e8d6125d652e2b24c71ef1e8ce4e011094f0.
Glossary and deeper reading · 8 refreshers
These editorial refreshers supplement the reading. Definitions and examples are added guidance; original source wording and recorded answers remain in their original context.
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Electronegativity
Electronegativity describes an atom's tendency to attract shared electron density in a bond. It helps predict bond polarity but is not the same as ionization energy or electron affinity. For example, oxygen attracts shared electrons more strongly than hydrogen in an O–H bond.
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Valence electron
Valence electrons are the electrons most directly involved in an atom's bonding and chemical behavior. For main-group atoms, they occupy the highest occupied principal shell. For example, neutral chlorine has seven valence electrons, helping explain its tendency to form a chloride ion.
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Electron shielding
Electron shielding is the reduction in nuclear attraction experienced by an electron because other electrons also repel it. Inner electrons generally shield an outer electron more effectively than electrons in the same shell. For example, adding an occupied inner shell helps explain why atomic size usually increases down a periodic-table group.
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Bond enthalpy
A bond enthalpy is the enthalpy required to break a specified bond in the gas phase. Tables often give average values, so estimates for complex reactions are approximate. For example, a reaction estimate adds the costs of bonds broken and subtracts the energy released by bonds formed.
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Covalent bond
A covalent bond involves shared electron density between atoms. The sharing can be approximately equal or unequal depending on the atoms and their environment. For example, the bond in a hydrogen molecule shares electrons between two identical atoms.
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Partial charge
A partial charge describes an uneven distribution of electron density within a bond or molecule. It is marked by a delta symbol and does not necessarily represent a full ionic charge. For example, oxygen carries a partial negative charge in water while its hydrogens carry partial positive charges.
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Ion
An ion is an atom or group of atoms with a net electric charge. Ordinary ion formation changes electron count without changing the element's proton count. For example, a sodium atom losing one electron becomes a sodium ion with charge +1.
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Electron delocalization
Electron delocalization means that an electronic description extends across multiple atoms rather than being confined to one localized bond or atom. It can affect bond equivalence and charge distribution. For example, nitrate's bonding electrons are described across several nitrogen–oxygen positions.