2.4 Structure of Metals and Alloys

What are metals and alloys in AP Chemistry?

Metallic solids are best modeled as positive metal ions sitting in a "sea" of delocalized valence electrons, which explains why metals conduct electricity, are malleable and ductile, and tend to have high melting points. Alloys mix a metal with other elements two main ways: interstitial alloys (small atoms squeeze into the gaps, like carbon in steel) and substitutional alloys (similar-sized atoms swap in, like zinc for copper in brass).

Why This Matters for the AP Chemistry Exam

This topic is about building and using models. You should be able to draw or describe a metallic solid as cations in a sea of delocalized electrons and use that picture to explain real properties like conductivity, malleability, and high melting points. AP Chemistry often asks you to connect what happens at the particle level to what you observe in the lab, and metals are a clean example of that.

You also need to tell apart the two alloy types and justify why one forms instead of the other based on atomic size. Expect to compare metals with ionic, molecular, and covalent-network solids, since questions like to test whether you can predict properties from structure.

Key Takeaways

• Model a metallic solid as positive metal ions surrounded by delocalized valence electrons (the "sea of electrons").
• Delocalized electrons explain metal properties: electrical conductivity, malleability, ductility, shiny appearance, and generally high melting points.
• Interstitial alloys form when much smaller atoms fill the spaces between larger metal atoms (carbon in iron makes steel).
• Substitutional alloys form when similar-sized atoms replace each other in the lattice (zinc replacing copper makes brass).
• Alloys are usually harder, stronger, and less malleable than pure metals because added atoms distort the lattice.
• Be ready to compare metallic solids with ionic, molecular, and covalent-network solids and predict properties from structure.

Metallic Bonding

Ionic and covalent substances are more common, but you also need to know metallic substances and their structure: a lattice of cations surrounded by a "sea" of valence electrons.

Sea of Electrons

When metal atoms form a metallic solid, their valence electrons become delocalized, leaving behind positive ions, or cations. Metallic bonding can be represented as an array of cations surrounded by this "sea" of valence electrons.

The nucleus and core electrons of each metal atom stay in place, but the valence electrons are very mobile, so we picture them as a shared sea. In most substances electrons belong to a specific atom, but in metals they move so freely that they aren't tied to any single atom.

Because the valence electrons move throughout the whole structure, metals have several distinctive properties:

• Good conductors of electricity: The delocalized valence electrons are mobile and move freely throughout the structure. This lets metals carry an electric current, which also connects to redox reactions in a later unit.
• High melting and boiling points: Metallic bonds are strong and take a lot of energy to break. Think about how much heat it takes to melt a metal like iron or gold.
• Shiny appearance: Light reflecting off the delocalized electrons gives many metals their shine.
• Malleability and ductility: Malleability is the ability to be hammered or shaped without breaking; ductility is the ability to be drawn into a wire. The metallic structure is less rigid than an ionic lattice, so the atoms can shift past each other without shattering.

Comparing Solids

When you compare properties across solid types, this chart is a useful reference:

For this topic, focus on the two bolded rows. Molecular and covalent-network solids are covered in more depth in Unit 3 when intermolecular forces come up.

Alloys

Metals can also combine with other elements to form alloys. An alloy forms when two or more elements, at least one of which is a metal, are melted and mixed together, then cooled so the solid mixture sets. The ratio of elements you mix in controls the properties of the final alloy, so each combination gives a material with its own characteristics.

There are two types of alloys you need to know.

Interstitial Alloys

Interstitial alloys form between atoms of significantly different radii, where the smaller atoms fill the interstitial spaces (the gaps) between the larger atoms.

The classic example is steel, which is iron with carbon. The carbon atoms are small enough to fit into the spaces in the iron lattice. The properties of the steel depend on the ratio of carbon to iron, so changing how much carbon you add produces different kinds of steel.

Interstitial alloys like steel tend to be strong and hard. The small atoms packed into the gaps make it harder for the layers of metal atoms to slide, which strengthens the material.

Substitutional Alloys

Substitutional alloys form between atoms of comparable radius, where one atom replaces another in the lattice. Unlike interstitial alloys, no extra atoms squeeze into gaps; instead, atoms swap places in the existing arrangement.

The common example is brass, where zinc atoms substitute for some of the copper atoms in the lattice. Because brass still has delocalized electrons, it keeps good electrical and thermal conductivity.

Interstitial vs. Substitutional Alloys

The key difference is what happens to the lattice:

• Interstitial: smaller atoms are added into the gaps between larger atoms (carbon in iron).
• Substitutional: similar-sized atoms replace atoms already in the lattice (zinc for copper).

In both cases the added atoms distort the regular metal structure. That distortion makes alloys generally harder and stronger, but usually less malleable, than the pure metals they come from.

How to Use This on the AP Chemistry Exam

Free Response

If a question asks you to model or describe a metallic solid, name the parts clearly: positive metal ions in a sea of delocalized valence electrons. Then connect that model to whatever property the question asks about. For conductivity, say the delocalized electrons are mobile and can carry charge. For malleability, say the layers of ions can shift without breaking the bonding because the electron sea holds them together.

Problem Solving

To decide between interstitial and substitutional alloys, compare atomic sizes. Significantly different radii point to an interstitial alloy (small atom fills gaps). Comparable radii point to a substitutional alloy (one atom replaces another). Use the steel and brass examples as your reference points.

Common Trap

When a question asks why a metal conducts but an ionic solid does not, the difference is mobile charge carriers. Metals have mobile delocalized electrons in the solid state. A solid ionic compound has its ions locked in a lattice, so it does not conduct until it is melted or dissolved, which frees the ions to move.

Common Misconceptions

• Metallic bonds are not the same as ionic bonds. An ionic solid has separate cations and anions held by electrostatic attraction. A metallic solid has only positive metal ions held together by shared, delocalized electrons.
• Delocalized electrons are not "free" to leave the metal. They move throughout the structure but are still attracted to the cations, which is why the solid holds together.
• Alloys are not pure compounds with fixed formulas. They are mixtures of elements in adjustable ratios, which is why you can make different grades of steel by changing the carbon amount.
• Interstitial does not mean "bigger atoms added." The atoms filling the gaps must be significantly smaller; if the atoms are similar in size, you get a substitutional alloy instead.
• A solid ionic compound does not conduct just because it contains ions. The ions have to be mobile, which means melting it or dissolving it in water.

Check Your Understanding

This practice question is based on one from the Advanced Placement YouTube channel and covers content from this guide and the previous guide.

(1) A student ran an experiment to see if the following solids conduct electricity.

(a) Explain the results the student saw.

(b) Is there anything that could have been different in this experiment to see the FeCl2 sample conduct electricity?

Practice Question Sample Responses

Sample responses for part (a):

• The sample of iron conducted electricity because it is a metal. Metals have delocalized valence electrons, shown by the sea of electrons model, which lets them conduct electricity.
• The sample of FeCl2 did not conduct electricity because it is an ionic solid. Its ions are locked in a lattice, so charges cannot move freely and the solid does not conduct.

For part (b):

As long as mobile charge carriers are present, the sample can conduct. Either response works:

• Melt the FeCl2 solid and then test it. Liquid FeCl2 conducts because the ions become mobile and can flow.
• Dissolve the FeCl2 in water. In solution the ions are free to move and carry charge.

Related AP Chemistry Guides

• 2.1 Types of Chemical Bonds
• Unit 2 Overview: Compound Structure and Properties
• 2.2 Intramolecular Force and Potential Energy
• 2.3 Structure of Ionic Solids
• 2.7 VSEPR and Bond Hybridization
• 2.6 Resonance and Formal Charge