AP Chemistry Unit 2 Notes: Chemical Bonding, Energy, and Solid Structure

Types of Chemical Bonds

A chemical bond is the attractive force that holds atoms or ions together in a stable structure. In AP Chemistry, you learn bonding not as a set of disconnected categories, but as a way to explain structure and properties—why some substances conduct electricity, why some are brittle solids, and why others are soft and malleable.

At a deep level, bonding is always about electrostatic attraction between positive and negative charges. What changes from bond to bond is (1) where the electrons are and (2) how the particles are arranged.

The three major bond types (and what they really mean)

Ionic bonding

Ionic bonding is the electrostatic attraction between cations (positive ions) and anions (negative ions). It commonly forms when one atom can lose electrons relatively easily (often a metal) and another can gain electrons relatively easily (often a nonmetal).

What’s easy to miss: ionic bonding is not a “one-ion-to-one-ion” connection like a line in a drawing. In a real ionic solid, each ion is attracted to many oppositely charged neighbors in a 3D ionic lattice. That network attraction is why ionic compounds tend to have high melting points.

Why it matters:

• Explains high melting/boiling points (strong, extended attractions)
• Explains electrical conductivity when molten or dissolved (mobile ions)
• Explains brittleness (lattice shifts can line up like charges)

Covalent bonding

A covalent bond forms when atoms share electron density. You can think of it as both nuclei being attracted to the same shared electrons.

Covalent bonding often forms between nonmetals and can produce:

• Molecules (discrete units like \text{H}_2\text{O} or \text{CO}_2)
• Network covalent solids (giant covalent lattices like diamond or \text{SiO}_2)

A key idea is that covalent bonding is not always “equal sharing.” If one atom attracts electrons more strongly, the bond becomes polar covalent, meaning electron density is pulled toward one side, creating partial charges.

Why it matters:

• Polarity affects solubility, molecular interactions, and sometimes reactivity
• Distinguishes molecular substances (often lower melting points) from network covalent solids (very high melting points)

Metallic bonding

Metallic bonding describes bonding in metals where positively charged metal cores (often described as cations in a lattice) are attracted to a “sea” of delocalized valence electrons.

This model explains the signature properties of metals:

• Electrical conductivity (electrons move through the solid)
• Malleability and ductility (layers can slide without breaking directional bonds)
• Luster (interaction of light with mobile electrons)

Bond type is a spectrum, not a set of boxes

A common misconception is that a compound is either 100% ionic or 100% covalent. In reality, bonding exists on a continuum:

• Large electronegativity differences often produce mostly ionic character.
• Smaller differences often produce mostly covalent character.

Electronegativity difference is a helpful guideline, but AP questions typically want you to justify using particle-level reasoning: Where is electron density? What charges exist? What structure forms? What property follows?

Connecting bonding to observable properties

Bonding type helps you predict macroscopic properties because structure determines how particles move and how strongly they attract.

Examples (with reasoning)

Example 1: Classify bonding and predict a property

• \text{NaCl}: metal + nonmetal, forms \text{Na}^+ and \text{Cl}^-, so an ionic lattice forms. Predict: high melting point; nonconductor as a solid; conductor when molten.
• \text{CO}_2: nonmetal + nonmetal, discrete molecules with covalent bonds. Predict: much lower melting/boiling points than ionic solids.

Example 2: Explaining metallic behavior

• Copper wire conducts because valence electrons are delocalized and can move under an applied electric field. If you tried to explain this with “ions moving,” you’d be describing molten ionic compounds instead—different mechanism.

Exam Focus

• Typical question patterns:
  • Given a formula (or element types), identify bonding type and predict properties like melting point or conductivity.
  • Compare substances and justify which has stronger attractions (ionic lattice vs molecular, metal vs ionic, etc.).
  • Explain a property at the particle level (why ionic solids are brittle; why metals conduct).
• Common mistakes:
  • Treating ionic bonding as a single “pair bond” rather than a lattice attraction to many neighbors.
  • Saying “ionic compounds conduct electricity” without specifying state (solid vs molten/aqueous).
  • Using electronegativity difference as the only justification instead of connecting to structure and charge mobility.

Intramolecular Force and Potential Energy

Bonding is fundamentally an energy story: stable structures form because the system can reach a lower potential energy state than the separated particles. In AP Chemistry, you’re expected to connect (1) electrostatic attractions/repulsions to (2) potential energy curves to (3) measurable quantities like bond energy and lattice energy.

What “intramolecular force” means

An intramolecular force is the force holding atoms together within a particle or solid structure (ionic attractions, covalent bonds, metallic bonding). These are generally much stronger than intermolecular forces (attractions between molecules), and that strength difference often explains why breaking chemical bonds requires significant energy.

Why potential energy depends on distance

Consider two particles that interact electrostatically (ions, or the nuclei and electrons in a covalent bond). As they approach, two competing effects occur:

• Attraction lowers potential energy (favorable).
• Repulsion (like-charge repulsion or electron cloud overlap) raises potential energy sharply at very short distances.

This creates a typical “well-shaped” potential energy curve vs. distance.

The key point on a potential energy curve: equilibrium bond length

The equilibrium bond length (or equilibrium distance between ions in an ionic pair model) is where potential energy is minimized. At this distance:

• The net force is zero (attraction balances repulsion).
• The system is most stable.

If particles get pulled farther apart than equilibrium, attraction dominates and pulls them back together. If pushed too close, repulsion dominates and pushes them apart.

Bond energy and what its sign means

The bond energy (often called bond dissociation enthalpy for covalent bonds) is the energy required to break one mole of a specific bond in the gas phase. In AP Chemistry, bond breaking is treated as endothermic (requires energy input), while bond formation is exothermic (releases energy).

If a bond dissociation enthalpy is reported as a positive number, that is consistent with “energy required to break.” The reverse process (forming the bond) would release the same magnitude of energy (conceptually), although real processes may involve multiple bonds and states.

Electrostatic models used for ionic interactions (Coulomb’s law)

For ionic compounds, the strength of attraction depends on charge magnitude and distance. A simplified model begins with Coulomb’s law for force:

F = k\frac{q_1 q_2}{r^2}

• F is electrostatic force
• k is a proportionality constant
• q_1 and q_2 are charges
• r is distance between charges

Potential energy for two point charges is proportional to:

U = k\frac{q_1 q_2}{r}

Interpretation you should be able to state clearly:

• Larger magnitude charges (like 2+ and 2-) make the attraction much stronger.
• Smaller ionic radii (smaller r) increase attraction.
• Opposite charges make q_1 q_2 negative, giving negative potential energy (a stable, lowered-energy state relative to infinite separation).

Real ionic solids are not just two ions, so AP often expects proportional reasoning rather than exact calculation: lattice energy increases with higher charges and smaller ionic sizes.

Lattice energy: translating “many attractions” into an energy trend

Lattice energy is the energy change associated with forming an ionic solid from gaseous ions (definitions may vary by sign convention in different texts). For AP-style reasoning, the crucial idea is: stronger ionic attractions correspond to a larger-magnitude lattice energy.

How to reason about trends without memorizing:

• Compare charge products: 1 \times 1 vs 2 \times 1 vs 2 \times 2
• Compare radii: smaller ions pack closer, reducing r

Worked trend examples (no plug-and-chug needed)

Example 1: Which has larger lattice energy, \text{NaCl} or \text{MgO}?

• \text{NaCl} uses \text{Na}^+ and \text{Cl}^- so the charge product magnitude is 1\cdot 1.
• \text{MgO} uses \text{Mg}^{2+} and \text{O}^{2-} so the charge product magnitude is 2\cdot 2 = 4.
• Also, \text{Mg}^{2+} is smaller than \text{Na}^+, which reduces distance.

Conclusion: \text{MgO} has significantly stronger attractions and thus a larger-magnitude lattice energy (and typically a higher melting point).

Example 2: Compare lattice energy of \text{LiF} vs \text{CsF}

• Charges are the same: +1 and -1.
• Radii differ: \text{Li}^+ is much smaller than \text{Cs}^+, so ions in \text{LiF} can get closer.

Conclusion: \text{LiF} has stronger attractions and larger-magnitude lattice energy.

What goes wrong: confusing strength with “number of bonds”

Students sometimes say an ionic compound is stronger because it has “more bonds.” That language is misleading. In an ionic lattice, each ion is attracted to multiple neighbors; strength is better explained by charge and distance rather than counting “bonds.”

Another common error is mixing up what bond energy refers to:

• Covalent bond energies usually refer to breaking a bond in a gas-phase molecule.
• Lattice energy refers to the ionic solid lattice (a bulk property).

Exam Focus

• Typical question patterns:
  • Compare bond strength or lattice energy using charge magnitude and ionic size (often framed as “which has higher melting point?”).
  • Interpret a potential energy vs. distance graph: identify equilibrium distance, bond energy (depth of well), and relative strength.
  • Explain why bond formation releases energy and bond breaking requires energy.
• Common mistakes:
  • Saying “stronger bond means longer bond length” (it’s generally the opposite for comparable bond types).
  • Forgetting that higher charges dramatically increase ionic attraction (a 2+ and 2- interaction is much stronger than 1+ and 1-).
  • Treating lattice energy as something you can infer from formula mass alone rather than electrostatics.

Structure of Ionic Solids

Ionic solids are not made of individual molecules. Instead, they form extended, repeating 3D arrangements called crystals. The structure is driven by a simple goal: arrange ions to maximize attractions between opposite charges while minimizing repulsions between like charges.

The ionic lattice: why the formula is an empirical ratio

In a compound like \text{NaCl}, the “1:1” formula does not mean one sodium is bonded to one chlorine as a pair. It means that in the entire crystal, the ratio of \text{Na}^+ to \text{Cl}^- is 1:1.

Each ion is surrounded by several oppositely charged ions. This local environment is described by the coordination number: how many nearest neighbors of opposite charge surround an ion.

Why it matters:

• Coordination and packing relate to density, stability, and lattice energy.
• Understanding “not molecular” prevents mistakes in predicting properties (like volatility or conductivity).

How ionic size influences structure

Ions cannot pack arbitrarily close because of electron cloud repulsion and the physical sizes of the ions. The relative sizes of cations and anions influence which crystal structures are stable.

A helpful qualitative rule: smaller cations can fit into “holes” between larger anions. If the cation is too small for a given hole, the structure becomes unstable and a different arrangement is favored.

You are not usually required to memorize hole geometry in AP Chemistry, but you should be able to explain that structure depends on both:

• electrostatic considerations (attraction/repulsion)
• size/packing considerations (how ions fit)

Properties explained by ionic structure

High melting points

To melt an ionic solid, you must disrupt a large number of strong ion-ion attractions throughout the lattice. That requires substantial energy, so ionic solids often melt at high temperatures.

Brittleness

Brittleness is one of the most testable lattice-structure consequences. When stress shifts one layer of ions relative to another, ions of the same charge can become aligned next to each other. Like charges repel strongly, and that repulsion can cause the crystal to cleave.

This is very different from metals, where delocalized electrons allow layers to slide without bringing like-charged ions into direct conflict.

Electrical conductivity depends on mobility

In a solid ionic lattice, ions are locked into position, so the solid does not conduct. When molten or dissolved in water, ions can move and carry charge, so conductivity increases.

Example explanations (particle-level)

Example 1: Why does solid \text{NaCl} not conduct electricity?
Even though it contains charged particles, those ions are fixed in a rigid lattice. Conductivity requires mobile charge carriers. In solid \text{NaCl}, charges cannot flow under an applied field, so it behaves as an insulator.

Example 2: Why does \text{NaCl} shatter when struck?
A blow can shift layers so that \text{Na}^+ ions are brought near other \text{Na}^+ ions (and similarly for \text{Cl}^-). The resulting repulsion is large, and the lattice fractures rather than deforming smoothly.

What goes wrong: treating ionic solids like molecular substances

A frequent misconception is to draw ionic compounds as discrete “molecules” and then try to apply molecular logic (like thinking they should be gases if “small”). The correct model is a bulk lattice whose properties depend on collective electrostatic forces.

Exam Focus

• Typical question patterns:
  • Explain brittleness vs malleability by describing what happens when layers shift.
  • Predict relative melting points using ionic charge and size (connecting to lattice energy).
  • Identify whether a formula represents a lattice (ionic solid) or discrete molecules and connect to properties.
• Common mistakes:
  • Claiming ionic solids conduct as solids because they “have ions” (mobility is the missing idea).
  • Describing the formula unit as a molecule rather than as a ratio in a lattice.
  • Explaining melting point using “strong bonds” but not specifying that the forces are extended ion-ion attractions throughout the crystal.

Structure of Metals and Alloys

Metals look simple on the surface, but their properties are among the clearest demonstrations that bonding is about electron behavior. Metallic bonding is best understood as a lattice of metal atoms (or metal cations) surrounded by valence electrons that are not tied to any one nucleus.

The electron-sea model (and what it explains)

In the electron-sea model, metal atoms contribute valence electrons to a shared pool of delocalized electrons. The positively charged metal cores are attracted to this mobile negative charge.

Why it matters: this single idea can explain multiple properties at once.

Electrical conductivity

Because electrons are delocalized, they can move through the solid when an electric potential is applied. This is fundamentally different from ionic conduction, which requires ions to move (typically only possible when molten or aqueous).

Thermal conductivity

Mobile electrons also transfer kinetic energy efficiently, helping explain why metals conduct heat well.

Malleability and ductility

Metals can deform without shattering because metallic bonding is non-directional. If a layer of metal atoms shifts, the delocalized electrons can redistribute and continue to hold the structure together. You don’t get the catastrophic like-charge alignment problem that occurs in ionic crystals.

Metallic crystal structure (qualitative)

Metals form crystalline lattices with repeating patterns. While AP Chemistry typically emphasizes properties and bonding models more than unit-cell memorization, it is still important to understand the big structural idea:

• Metal atoms pack closely in an ordered array.
• Delocalized electrons occupy the space between and around atoms.

This close packing is consistent with metals often having relatively high densities.

Alloys: tuning properties by mixing atoms

An alloy is a solid mixture of elements where at least one is a metal. Alloying changes properties because it disrupts the regularity of the metal lattice and alters how easily layers slide.

There are two major structural categories you should be able to describe.

Substitutional alloys

In a substitutional alloy, atoms of a second element replace some of the metal atoms in the lattice.

How it changes properties:

• If the substituting atoms are similar in size, they can fit into the lattice.
• The lattice becomes less uniform, often making it harder for layers to slide perfectly.

A classic example is brass (primarily copper and zinc), where zinc atoms substitute into a copper lattice.

Interstitial alloys

In an interstitial alloy, smaller atoms fit into the holes (interstices) between metal atoms.

How it changes properties:

• Interstitial atoms can strongly hinder the movement of metal layers.
• This often increases hardness and strength.

Steel is a common example: small carbon atoms occupy spaces within an iron lattice (along with other possible alloying elements depending on the steel).

Connecting metallic structure to macroscopic behavior

Example 1: Why is pure copper relatively soft compared with many alloys?
In a relatively uniform lattice, layers can slide more easily under force. When you add alloying atoms, you distort the lattice and create obstacles to sliding, making deformation more difficult.

Example 2: Why do metals remain conductive when alloyed (often less, but still conductive)?
Delocalized electrons still exist in alloys, so charge carriers remain available. However, lattice disruptions can scatter electrons more, sometimes increasing electrical resistance.

What goes wrong: confusing metallic bonding with ionic bonding

Because the electron-sea model sometimes describes “metal cations in a sea of electrons,” students may incorrectly conclude that metallic solids should be brittle like ionic solids. The key difference is that the negative charge is mobile and spread out, so when layers shift, the attractive interactions can re-form rather than snapping the structure apart.

Exam Focus

• Typical question patterns:
  • Explain conductivity and malleability of metals using delocalized electrons.
  • Compare metals vs ionic solids: why metals deform while ionic solids fracture.
  • Describe how alloys (substitutional vs interstitial) change hardness/strength in terms of lattice disruption.
• Common mistakes:
  • Saying metals conduct because “ions move” in the solid (it’s electrons in the solid; ions are fixed).
  • Claiming alloys always lower melting point or always increase strength (property changes depend on composition and structure).
  • Explaining malleability without mentioning non-directional bonding and electron mobility.