This lab is really asking one question: can you look at how a solid behaves and figure out what's holding it together? You're collecting macroscopic evidence (things you can observe and measure) and using it to make claims about the particulate-level structure of unknown substances. That's the core skill here, and it shows up all over the AP exam.

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Why This Lab Matters for the AP Exam

The AP Chemistry exam regularly asks you to connect observable properties to bonding type. You might get a free-response question that gives you melting point data, conductivity results, or a description of hardness and asks you to identify the solid type and justify your answer. This lab is literally practice for that exact move.

You also need to be comfortable drawing or interpreting particulate models of solids. The exam will ask you to represent ionic, metallic, molecular, and covalent network solids at the particle level and explain why each type behaves the way it does. This lab builds that reasoning from the ground up.

CED Connections

This lab pulls from three topics across Units 2 and 3.

Topic 2.3: Structure of Ionic Solids

Learning Objective 2.3.A: Represent an ionic solid with a particulate model consistent with Coulomb's law and the properties of the constituent ions.
Essential Knowledge 2.3.A.1: Cations and anions arrange in a 3-D array that maximizes attraction and minimizes repulsion.

Topic 2.4: Structure of Metals and Alloys

Learning Objective 2.4.A: Represent a metallic solid and/or alloy using a model that shows essential characteristics of the structure and interactions.
Essential Knowledge 2.4.A.1: Metallic bonding is represented as positive metal ions in a "sea of electrons" (delocalized valence electrons).
Essential Knowledge 2.4.A.2 and 2.4.A.3: Interstitial alloys (like steel) and substitutional alloys (like brass) differ based on atomic radius relationships.

Topic 3.2: Properties of Solids

Learning Objective 3.2.A: Explain the relationship among macroscopic properties, particulate-level structure, and interparticle interactions.
Essential Knowledge 3.2.A.1 through 3.2.A.6: Each solid type (ionic, covalent network, molecular, metallic) has a distinct set of properties that follow directly from its bonding and structure.

What You Need to Be Able to Do

This lab builds several skills that show up directly on the AP exam.

Identify solid type from property data. Given conductivity, melting point, hardness, or solubility results, you should be able to classify a solid as ionic, metallic, molecular, or covalent network.
Justify claims with evidence and reasoning. This is the classic claim-evidence-reasoning (CER) format. Your claim is the solid type. Your evidence is the data. Your reasoning connects the data to the bonding model.
Draw or interpret particulate models. You should be able to sketch what each solid type looks like at the particle level and label the relevant interactions.
Apply Coulomb's law qualitatively. You should be able to explain how charge magnitude and interionic distance affect the strength of attraction in an ionic solid.
Compare properties across solid types. Ranking melting points or explaining why one solid conducts and another doesn't requires you to reason about interaction strength.
Design or evaluate an investigation. You should understand what each test is measuring and why it's useful as evidence.

Core Concepts

The Four Solid Types

Every solid you encounter in this lab falls into one of four categories. Each has a different structure at the particle level, and that structure explains everything about how it behaves.

Ionic solids are made of cations (positively charged ions) and anions (negatively charged ions) arranged in a repeating 3-D structure called a crystal lattice. The force holding them together is electrostatic force, which is the attraction between opposite charges. Because these forces are strong, ionic solids have high melting points and high boiling points. They're also brittle: if you shift one layer of ions, like charges end up next to each other and the repulsion causes the crystal to crack. Ionic solids only conduct electricity when the ions can move, meaning when the solid is dissolved in water or melted into a liquid.

Metallic solids are made of positive metal ion cores arranged in a lattice, surrounded by delocalized valence electrons that move freely throughout the structure. This "sea of electrons" model explains why metals conduct electricity and heat so well. It also explains malleability (the ability to be hammered into sheets) and ductility (the ability to be drawn into wires): the electron sea allows metal cores to slide past each other without breaking bonds.

Covalent network solids are made of atoms connected by covalent bonds that extend throughout the entire solid in a continuous network. Diamond is the classic example: every carbon atom is covalently bonded to four others in a rigid 3-D structure. This makes diamond extremely hard and gives it a very high melting point. Graphite is also a covalent network solid, but its structure is layered 2-D sheets. The layers can slide past each other, which is why graphite is soft and slippery. Covalent network solids do not conduct electricity (with the exception of graphite, which conducts along its layers due to delocalized electrons within each sheet).

Molecular solids are made of individual covalently bonded molecules held together by intermolecular forces (IMFs) like London dispersion forces, dipole-dipole interactions, or hydrogen bonding. Because IMFs are much weaker than ionic, metallic, or covalent bonds, molecular solids have low melting points. They do not conduct electricity because the valence electrons are locked inside the covalent bonds of each molecule.

Coulomb's Law and Ionic Solids

Coulomb's law describes the force between two charged particles. The key relationship for AP Chemistry is:

$F \propto \frac{q_1 \cdot q_2}{r^2}$

Where $q_1$ and $q_2$ are the charges on the ions and $r$ is the distance between them (the interionic distance). A larger charge magnitude means a stronger attraction. A smaller interionic distance (smaller ionic radius) also means a stronger attraction. This directly affects lattice energy, which is the energy released when gaseous ions come together to form an ionic solid. Higher lattice energy means a more stable, harder-to-melt ionic solid.

So MgO (charges of 2+ and 2-) has a much higher melting point than NaCl (charges of 1+ and 1-) because the electrostatic attraction is stronger. That's Coulomb's law in action.

Lattice Energy

Lattice energy (sometimes called lattice enthalpy) is the energy required to completely separate one mole of an ionic solid into its gaseous ions. It's a measure of how strongly the ions are held together. Higher lattice energy means higher melting point, lower solubility in many cases, and a harder crystal. Lattice energy increases when ion charges are higher or when ionic radii are smaller, both of which decrease interionic distance and increase electrostatic attraction.

How the Lab Works

The investigation gives you a set of unknown solid samples, each in a labeled bottle. Your job is to run a series of property tests and use the results to figure out what type of solid each one is.

The tests you run are chosen because each one probes a different aspect of bonding. Here's the logic behind each type of test:

Conductivity testing checks whether a solid can carry an electric current. You test the solid in its pure form, dissolved in water, and sometimes when melted. An ionic solid won't conduct as a pure solid (ions are locked in place) but will conduct when dissolved or melted (ions become mobile). A metallic solid conducts in all states. A molecular solid won't conduct in any state. A covalent network solid generally won't conduct either (except graphite).

Melting point observation tells you about the strength of the interparticle interactions. A substance that melts at a very high temperature has strong interactions holding its particles together. A substance that melts at a low temperature has weak interactions. This helps you distinguish between, say, a molecular solid (low melting point) and an ionic or covalent network solid (high melting point).

Physical property observations like hardness, brittleness, and whether the solid can be deformed give you more clues. A solid that shatters when struck is likely ionic. A solid that bends or flattens is likely metallic. A solid that is extremely hard and doesn't deform is likely a covalent network solid.

Solubility in water can also help. Many ionic solids dissolve in water because water molecules can stabilize the separated ions. Molecular solids may or may not dissolve depending on polarity. Metals and covalent network solids generally do not dissolve in water.

The key move in this lab is triangulating. No single test tells you everything. You combine all your observations to build a case for one solid type.

Data and Analysis Moves

Organizing Your Data

A comparison table is your best friend here. Set up columns for each property you tested (conductivity as solid, conductivity in solution, melting point range, hardness, solubility) and rows for each unknown sample. Once the table is filled in, patterns become much easier to see.

Making the Classification

Use the property patterns to match each unknown to a solid type. Here's a quick reference for what each type looks like in your data:

Property	Ionic	Metallic	Molecular	Covalent Network
Conducts as solid	No	Yes	No	No (except graphite)
Conducts in solution	Yes	N/A	No	No
Melting point	High	Variable, often high	Low	Very high
Hardness	Hard, brittle	Malleable/ductile	Soft	Very hard (3-D) or soft layers (graphite)

Writing Your Justification

For each unknown, your analysis should follow this structure:

Claim: State the solid type.
Evidence: List the specific data values or observations from your tests.
Reasoning: Explain why that data is consistent with the bonding model. Reference the particle-level structure. For example, if a solid conducts only in solution, explain that ionic solids require mobile ions to conduct, and dissolving releases those ions from the lattice.

Comparing Ionic Solids

If you have two ionic solids in the lab, you might be asked to compare their melting points or lattice energies. Go back to Coulomb's law. Look at the charges of the ions and the sizes of the ions (ionic radius). Higher charges and smaller radii mean stronger electrostatic attraction, higher lattice energy, and higher melting point.

Particulate Models

You may be asked to draw a particulate model for one or more of your unknowns. For an ionic solid, draw alternating cations and anions in a regular array. For a metallic solid, draw positive ion cores with dots representing delocalized electrons spread throughout. For a molecular solid, draw individual molecules with dashed lines or labels indicating IMFs between them. For a covalent network solid, draw atoms connected by lines representing covalent bonds extending in all directions.

Common Mistakes

Thinking ionic solids always conduct. They only conduct when ions can move. A solid chunk of NaCl sitting on a table does not conduct. This trips up a lot of students on the exam.

Confusing intermolecular forces with intramolecular bonds. In a molecular solid, the covalent bonds inside each molecule are strong. But the forces between molecules are weak IMFs. Melting a molecular solid breaks the IMFs, not the covalent bonds. The molecules stay intact.

Assuming high melting point always means ionic. Covalent network solids like diamond and silicon dioxide have extremely high melting points too. You need conductivity and other data to tell them apart from ionic solids.

Forgetting that graphite is the exception. Graphite is a covalent network solid that conducts electricity along its layers because of delocalized electrons within each 2-D sheet. If you see a solid that is soft, dark, and conducts, don't automatically call it a metal.

Mixing up malleability and ductility. Malleability is the ability to be hammered into flat sheets. Ductility is the ability to be drawn into wires. Both apply to metals and both come from the same source: the electron sea allows metal cores to rearrange without breaking bonds.

Applying Coulomb's law backwards. Larger ionic radius means greater interionic distance, which means weaker attraction and lower lattice energy. Students sometimes flip this relationship.

Saying a substance "has no intermolecular forces." Every substance has at least London dispersion forces. The question is always about the type and strength of IMFs, not whether they exist.

Quick Review Checklist

You can classify a solid as ionic, metallic, molecular, or covalent network based on conductivity, melting point, hardness, and solubility data.
You can explain why ionic solids conduct only when dissolved or melted, using the idea that ions must be mobile to carry charge.
You can apply Coulomb's law qualitatively: higher ion charges and smaller ionic radii both increase lattice energy and melting point.
You can draw a particulate model for each solid type that shows the correct particles and the interactions between them.
You can distinguish between the 3-D network structure of diamond (hard, nonconducting) and the layered structure of graphite (soft, conducting along layers).
You can write a complete CER response that connects macroscopic property data to a specific bonding model.
You know that metallic malleability and ductility come from the electron sea model, and that adding interstitial atoms (like carbon in steel) reduces those properties by making the lattice more rigid.