Prepare for AP Chemistry Topic 2.1 with this free practice quiz. Test your understanding of ionic, covalent, and metallic bonds, and how electronegativity differences determine bond character to maximize your AP score.
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The Pre-Quiz Review
1. Ionic Bonding
An ionic bond is the electrostatic attraction between oppositely charged ions, usually formed when a metal transfers electrons to a nonmetal.
- Metals tend to lose electrons and form cations.
- Nonmetals tend to gain electrons and form anions.
- The resulting ions attract because of Coulombic forces.
- Ionic compounds form crystal lattices, not individual molecules.
Formation of ions
Main-group elements commonly form ions that achieve a noble-gas-like electron configuration:
- Group 1 → +1
- Group 2 → +2
- Group 13 → +3
- Group 15 → −3
- Group 16 → −2
- Group 17 → −1
Examples:
- Na → Na⁺ + e⁻
- Mg → Mg²⁺ + 2e⁻
- Cl + e⁻ → Cl⁻
- O + 2e⁻ → O²⁻
For an ionic compound, total positive charge must equal total negative charge.
Examples:
- Na⁺ + Cl⁻ → NaCl
- Mg²⁺ + 2Cl⁻ → MgCl₂
- 2Al³⁺ + 3O²⁻ → Al₂O₃
A common MCQ trap is confusing the subscripts in a chemical formula with charges. In MgCl₂, the ₂ means two chloride ions are present; it does not mean Cl has a −2 charge.
Ionic compounds and properties
Ionic compounds generally:
- Have high melting and boiling points.
- Are crystalline solids at room temperature.
- Are brittle.
- Conduct electricity when molten or dissolved in water because ions can move.
- Usually do not conduct electricity as solids because the ions are fixed in the lattice.
The strength of ionic attraction increases as:
- Ion charges become larger.
- Ionic radii become smaller.
Thus, MgO has stronger electrostatic attractions than NaCl because Mg²⁺/O²⁻ have larger charge magnitudes than Na⁺/Cl⁻.
This connects to lattice energy: stronger ionic attractions generally correspond to larger-magnitude lattice energies and higher melting points.
2. Covalent Bonding
A covalent bond forms when atoms share electrons. Covalent bonding is especially common between nonmetals.
The shared electrons are attracted to the nuclei of both bonded atoms.
Single, double, and triple bonds
- Single bond: 1 shared electron pair
- Double bond: 2 shared electron pairs
- Triple bond: 3 shared electron pairs
In general:
Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.
Therefore:
- Bond strength: triple > double > single
- Bond length: single > double > triple
This relationship is frequently tested.
Lewis structures
For Lewis structures:
- Count total valence electrons.
- Choose a reasonable skeletal structure.
- Connect atoms with single bonds.
- Complete terminal atoms’ octets.
- Place remaining electrons on the central atom.
- Form multiple bonds if necessary to satisfy octets and reduce formal charges.
Important exceptions:
- H follows a duet, not an octet.
- B and Be can have fewer than eight electrons in some stable compounds.
- Elements in period 3 and beyond can sometimes have expanded valence shells.
For ions, add electrons for negative charge and remove electrons for positive charge.
Example: SO₄²⁻ has two more electrons than the neutral atoms would provide.
Formal charge
Formal charge helps determine the most reasonable Lewis structure:
Formal charge = valence electrons − nonbonding electrons − number of bonding electrons/2
Prefer structures that generally:
- Minimize formal charges.
- Place negative formal charge on the more electronegative atom when possible.
- Have the correct total charge.
A structure with formal charges is not automatically incorrect. Resonance structures can distribute charge among equivalent atoms.
3. Electronegativity
Electronegativity measures an atom’s tendency to attract shared electrons in a chemical bond.
Periodic trend:
- Increases from left to right across a period.
- Increases from bottom to top within a group.
- Fluorine has the highest electronegativity.
A useful ranking:
F > O > Cl > N > C > H approximately.
Do not confuse electronegativity with electron affinity or ionization energy.
- Electronegativity: attraction for shared electrons in a bond.
- Ionization energy: energy required to remove an electron from a gaseous atom.
- Electron affinity: energy change associated with adding an electron to a gaseous atom.
4. Electronegativity Differences and Bond Type
The difference in electronegativity between bonded atoms helps describe bond polarity.
Calculate the difference conceptually as:
ΔEN = |EN₁ − EN₂|
The larger the difference, the greater the unequal sharing of electrons.
General classification:
- Very small difference → nonpolar covalent
- Intermediate difference → polar covalent
- Very large difference → bonding has substantial ionic character
Do not treat cutoff values as universal laws. Different textbooks use slightly different numerical boundaries. AP Chemistry questions generally emphasize the underlying trend rather than requiring a particular arbitrary cutoff.
Nonpolar covalent bonds
Electrons are shared approximately equally.
Examples:
- H–H
- Cl–Cl
- O=O
A bond between identical atoms has ΔEN = 0 and is nonpolar.
Polar covalent bonds
Electrons are shared unequally because one atom attracts the shared electrons more strongly.
Examples:
- H–Cl
- O–H
- C–O
The more electronegative atom acquires a partial negative charge (δ−), while the less electronegative atom acquires a partial positive charge (δ+).
For H–Cl:
Hδ+—Clδ−
The bond polarity points toward the more electronegative atom.
5. Bond Polarity vs. Molecular Polarity
This distinction is one of the most important MCQ concepts.
A molecule can contain polar bonds but be nonpolar overall.
Molecular polarity depends on both:
- Bond polarities
- Molecular geometry
Example: CO₂ contains two polar C=O bonds, but the molecule is linear:
O=C=O
The two bond dipoles point in opposite directions and cancel. Therefore, CO₂ is nonpolar overall.
By contrast, H₂O has polar O–H bonds and a bent geometry. The bond dipoles do not cancel, so H₂O is polar.
Common examples:
- CO₂ → polar bonds, nonpolar molecule
- H₂O → polar bonds, polar molecule
- CCl₄ → polar bonds, nonpolar molecule
- NH₃ → polar bonds, polar molecule
- CH₂Cl₂ → polar bonds, polar molecule
For MCQs, never determine molecular polarity from electronegativity alone. Draw or visualize the molecular geometry and determine whether the bond dipoles cancel.
6. Metallic Bonding
A metallic bond is the attraction between positively charged metal cores and a delocalized collection of valence electrons.
Metal atoms do not form discrete molecules in a metallic solid. Their valence electrons are relatively mobile throughout the structure.
This explains characteristic metallic properties:
- High electrical conductivity
- High thermal conductivity
- Malleability
- Ductility
- Metallic luster
Why metals conduct electricity
Metallic solids contain mobile, delocalized electrons. When an electric field is applied, these electrons can move through the material.
In contrast:
- Ionic solids have ions locked into place.
- Molecular solids generally lack mobile charged particles.
- Covalent network solids generally do not have freely moving electrons, with important exceptions such as graphite.
7. Comparing the Three Bonding Models
| Bonding type | Typical elements | Particles/structure | Main attraction | Electrical conductivity |
|---|---|---|---|---|
| Ionic | Metal + nonmetal | Crystal lattice of ions | Cation–anion electrostatic attraction | Conducts when molten/aqueous |
| Covalent | Nonmetal + nonmetal | Molecules or networks | Shared electrons attracted to nuclei | Usually poor |
| Metallic | Metals | Metal atoms/cores with delocalized electrons | Metal cores + delocalized electrons | Good |
Remember that covalent bonding can produce either molecular substances or network covalent solids.
Examples of network covalent substances include:
- Diamond
- Graphite
- SiO₂
These are not molecular solids made of separate SiO₂ molecules.
8. Bond Polarity: High-Yield MCQ Rules
When comparing two bonds:
- Larger ΔEN → generally greater bond polarity.
- The more electronegative atom carries δ−.
- The less electronegative atom carries δ+.
- Bond dipoles point toward the more electronegative atom.
When comparing compounds, do not assume the compound with the largest number of polar bonds is automatically the most polar molecule.
Geometry controls dipole cancellation.
Symmetry is especially important:
- Highly symmetrical molecules can have complete dipole cancellation.
- Asymmetrical molecules are more likely to have a net dipole.
9. Common AP Chemistry Traps
Ionic vs. polar covalent
A bond does not suddenly become “ionic” simply because its electronegativity difference is large. Bonding exists on a continuum from nonpolar covalent to polar covalent to strongly ionic character.
Partial vs. full charges
Polar covalent bonds produce partial charges (δ+ and δ−).
Ionic compounds contain ions with full integer charges, such as Na⁺ and Cl⁻.
Bond polarity vs. molecular polarity
A molecule may contain several polar bonds and still be nonpolar if the dipoles cancel.
Conductivity
Do not say that all ionic substances conduct electricity.
- Solid NaCl → poor conductor
- Molten NaCl → conducts
- NaCl(aq) → conducts
The key is whether charged particles can move.
Bond length vs. bond strength
Shorter covalent bonds are generally stronger.
For the same two atoms:
triple > double > single in bond strength.
Electronegativity trend
Moving up and right on the periodic table generally increases electronegativity.
Fluorine is the highest.
Metallic bonding
Metallic bonding is not electron transfer from one specific metal atom to one specific neighboring atom. The valence electrons are delocalized throughout the metallic structure.
10. Rapid MCQ Checklist
Before answering a bonding question, ask:
- What elements are involved?
- Metal + nonmetal → likely ionic.
- Nonmetal + nonmetal → likely covalent.
- Metal + metal → metallic bonding.
- Which atom is more electronegative?
- Electrons are pulled toward that atom.
- What is the ΔEN?
- Larger difference → greater bond polarity.
- Is the question asking about a bond or the whole molecule?
- Bond polarity depends mainly on ΔEN.
- Molecular polarity depends on bond dipoles and geometry.
- Can the dipoles cancel?
- Check symmetry and molecular shape.
- Are electrons or ions mobile?
- Mobile electrons → metallic conductivity.
- Mobile ions → molten/aqueous ionic conductivity.
- Are you comparing bond lengths or strengths?
- Stronger bond → generally shorter bond.
- Is the substance ionic, molecular, or network covalent?
- This determines many physical properties.
11. Must-Know Takeaways
- Ionic bonding: electrostatic attraction between oppositely charged ions.
- Covalent bonding: sharing of electron pairs between atoms.
- Metallic bonding: metal cores attracted to delocalized valence electrons.
- Electronegativity increases up and to the right.
- Fluorine is the most electronegative element.
- Greater ΔEN generally means greater bond polarity.
- More electronegative atom → δ−.
- Less electronegative atom → δ+.
- Polar bonds do not necessarily make a polar molecule.
- Molecular geometry determines whether bond dipoles cancel.
- Ionic solids generally do not conduct; molten or dissolved ionic compounds do.
- Metals conduct because of mobile, delocalized electrons.
- Triple bonds are generally shorter and stronger than double bonds; double bonds are shorter and stronger than single bonds.
- Always distinguish bond polarity from molecular polarity.