Structure of Metals and Alloys Study Notes- AP Chemistry- Unit 2.4 -New Syllabus 2024-2025

Structure of Metals and Alloys Study Notes- AP Chemistry

Structure of Metals and Alloys Study Notes as per latest AP Chemistry Syllabus

LEARNING OBJECTIVE

Represent a metallic solid and/or alloy using a model to show essential characteristics of the structure and interactions present in the substance.

Key Concepts:

Metallic Bonding: Sea of Delocalized Electrons
Interstitial Alloys: Smaller Atoms in Larger Atomic Spaces
Substitutional Alloys: Atoms of Comparable Radius

AP Chemistry-Concise Summary Notes- All Topics

2.4.A.1 Metallic Bonding: Sea of Delocalized Electrons:

1. Structure of Metals: The structure of metals is closely related to the arrangement of metal ions and the close packing of atoms in metallic solids.

i. Arrangement of Metal Ions:

In a metallic solid, metal atoms lose their valence electrons and form positively charged metal ions (cations).
These metal ions are arranged in a highly ordered pattern, creating a regular structure.
The valence electrons that are released form a “sea of electrons” that is free to move through the entire metal lattice. These delocalized electrons act as a glue, holding the metal ions together and allowing for the characteristics of metallic bonding (such as conductivity and malleability).

The arrangement of metal ions can be described using different types of crystal lattices. These lattices represent the regular, repeating patterns in which metal ions are packed together.

ii. Close Packing of Atoms in Metallic Solids:

Close packing refers to the most efficient arrangement of atoms within a material, where atoms are packed as closely as possible to minimize empty space.
In metallic solids, atoms arrange themselves in a close-packed structure to achieve maximum packing efficiency and stability.

There are two main types of close packing in metals:

a. Face-Centered Cubic (FCC) Structure:

In the FCC structure, each unit cell contains atoms at each corner and the center of each face of the cube. This creates a highly efficient packing arrangement.
The atoms in the FCC arrangement are packed closely, with a coordination number of 12 (each atom is surrounded by 12 nearest neighbors).
The packing efficiency of FCC is 74%, meaning 74% of the volume is filled with metal atoms, and the rest is empty space.
Metals that adopt an FCC structure include aluminum (Al), copper (Cu), gold (Au), and silver (Ag).

b. Hexagonal Close-Packed (HCP) Structure:

In the HCP structure, the metal ions are arranged in a hexagonal shape, with two layers of atoms forming a close-packed structure.
The atoms are arranged so that each atom has a coordination number of 12.
Like the FCC structure, HCP has a packing efficiency of 74%.
Some examples of metals that adopt the HCP structure include magnesium (Mg), titanium (Ti), and zinc (Zn).

3. Body-Centered Cubic (BCC) Structure (Less close-packed):

The BCC structure is less efficient in terms of packing compared to FCC and HCP.
In a BCC arrangement, there is an atom at each corner of the cube and one atom at the center of the cube.
The coordination number of BCC is 8 (each atom is surrounded by 8 nearest neighbors).
The packing efficiency of BCC is lower, at around 68%.
Examples of metals that form BCC structures include iron (Fe) at room temperature, chromium (Cr), and tungsten (W).

4. Importance of Close Packing in Metals:

Strength: The close-packed arrangement allows the metal to withstand greater forces without breaking because the atoms are packed tightly together.
Malleability and Ductility: The close-packed structures (FCC and HCP) also make metals more malleable and ductile, meaning they can be shaped and stretched without breaking. This is due to the ease with which layers of atoms can slip past each other when force is applied.
Density: Metals with close-packing structures generally have higher densities since the atoms are packed tightly together, leaving less empty space.

Summary of Close-Packed Structures:

FCC: 12 neighbors, 74% packing efficiency, found in metals like Cu, Al, Au.
HCP: 12 neighbors, 74% packing efficiency, found in metals like Ti, Mg, Zn.
BCC: 8 neighbors, 68% packing efficiency, found in metals like Fe, Cr, W.

In essence, the structure of metals and the arrangement of their metal ions play a crucial role in determining their physical properties. Close packing allows for the maximization of bonding strength, density, and ductility, while the metallic bond (the “sea of electrons”) enables the flexibility and conductivity that are typical of metals.

2. Delocalized Electrons:

Delocalized electrons refer to electrons in a metal that are not bound to any specific atom or ion. These electrons are free to move throughout the entire metal structure, which is a key characteristic of metallic bonding and is responsible for many of the properties of metals, such as electrical conductivity, thermal conductivity, and malleability.

i. Definition of valence electrons: Valence electrons are the outermost electrons of an atom that are involved in chemical bonding. These electrons are in the outermost shell of an atom and have the highest energy level. Valence electrons are crucial because they participate in the formation of bonds with other atoms.

In metals, the number of valence electrons typically ranges from 1 to 3, and these electrons are loosely bound to their respective atoms. This is why metals can easily lose these electrons and form positive metal ions.

ii. How Electrons Become Delocalized in Metals: In metallic solids, the metallic bond is formed by the electrostatic attraction between the positively charged metal ions (cations) and the “sea of electrons.”

Electron Loss and Formation of Metal Ions:
- In a metal, the valence electrons are not tightly bound to individual atoms because the difference between the energy levels of the valence shell and the inner shells is relatively small.
- When metals form a solid, the metal atoms lose their valence electrons, which are then freed from any one particular atom.
Sea of Electrons:
- The lost valence electrons do not remain with individual atoms but instead form a “sea of electrons” that are free to move throughout the structure.
- These delocalized electrons are free-moving and are not confined to a single atom or bond but instead flow through the entire lattice of metal ions.
Metallic Bonding:
- The metal cations (positively charged metal ions) are surrounded by this electron sea.
- The attraction between the metal cations and the delocalized electrons forms the metallic bond, which holds the metal atoms together in a lattice structure.
- The delocalized electrons act like a “glue” that binds the metal ions and also allows them to slide past each other without breaking the metallic bond, giving metals their malleability and ductility.
Movement of Delocalized Electrons:
- The delocalized electrons can move freely in response to electric fields, which is why metals conduct electricity.
- They can also transfer thermal energy, making metals good conductors of heat.

iii. Importance of Delocalized Electrons:

Electrical Conductivity: Because electrons are free to move, metals can conduct electricity. When an electric potential is applied, the electrons flow towards the positive side, carrying charge.
Thermal Conductivity: Delocalized electrons can also transfer kinetic energy, which helps metals conduct heat.
Malleability and Ductility: The ability of the electron sea to move allows metal ions to shift position without breaking the bond, making metals malleable (able to be hammered into sheets) and ductile (able to be stretched into wires).

3. Properties of Metals: Metals are known for their distinctive physical properties, which make them useful in various applications. The most important properties of metals include electrical conductivity, malleability and ductility, and thermal conductivity. These properties are largely due to the unique structure of metals and the presence of delocalized electrons in the metallic bond.

Property	Explanation	Examples
Electrical Conductivity	Delocalized electrons in metals allow them to carry electrical current.	Copper (Cu), Silver (Ag)
Malleability and Ductility	Delocalized electrons allow metal ions to slide past each other, making metals easy to shape or stretch.	Gold (Au), Aluminum (Al), Copper (Cu)
Thermal Conductivity	Free electrons in metals can transfer kinetic energy quickly, allowing heat to pass through them efficiently.	Copper (Cu), Aluminum (Al), Gold (Au)

The delocalized electrons in the metallic bond are crucial in giving metals their remarkable properties of electrical conductivity, thermal conductivity, and malleability/ductility. These properties make metals ideal for a wide range of industrial and technological applications, from electrical wiring to building materials to cooking utensils.

4. Nature of Metallic Bonding:

Metallic bonding is the type of chemical bonding found in metals. It explains the unique properties of metals such as electrical conductivity, malleability, ductility, and luster. The key feature of metallic bonding is the electrostatic attraction between metal ions and delocalized electrons, which leads to the formation of a cohesive structure that provides the metal with its distinct characteristics.

i. Electrostatic attraction between metal ions and delocalized electrons:

Metal Ions: In a metal, the outermost electrons (valence electrons) are not bound tightly to individual metal atoms. Instead, they are freed from their respective atoms and become delocalized electrons.
Delocalized Electrons: These electrons are free to move throughout the entire metal structure, creating a “sea of electrons”. The delocalized electrons are spread over many atoms and act as a “glue” holding the metal ions together.
Metal Ions: The metal atoms, having lost their outermost electrons, become positively charged metal ions (cations). These cations are arranged in a regular, repeating pattern known as a lattice.
Electrostatic Attraction: The key to metallic bonding is the electrostatic attraction between the positively charged metal ions (cations) and the negatively charged delocalized electrons. This attractive force holds the metal ions together in the lattice structure.
- The delocalized electrons are free to move, but they are still attracted to the metal ions. This attraction forms a stable bond throughout the metal.
Characteristics of Metallic Bonding:
- Conductivity: The delocalized electrons can move freely, allowing metals to conduct both electricity and heat.
- Malleability and Ductility: The metal ions can slide past each other due to the mobility of the delocalized electrons, which allows metals to be bent or stretched without breaking.

ii. Strength of metallic bonds:

The strength of metallic bonds depends on several factors, including the number of delocalized electrons, the size of the metal ions, and the arrangement of metal ions in the lattice.

i. Number of Delocalized Electrons:

Metals with more delocalized electrons tend to have stronger metallic bonds. This is because more electrons contribute to the “electron sea,” increasing the attraction between the electrons and metal ions, making the bond stronger.
For example:
- Alkali metals (e.g., sodium, Na) have fewer valence electrons, leading to weaker metallic bonds.
- Transition metals (e.g., copper, Cu, or gold, Au) have more valence electrons, resulting in stronger metallic bonds.

ii. Size of Metal Ions:

Smaller metal ions tend to form stronger metallic bonds. This is because smaller metal ions have a higher charge density (more positive charge per unit volume), and the delocalized electrons are more tightly attracted to them.
In smaller metal ions, the distance between the metal ions and delocalized electrons is shorter, resulting in a stronger electrostatic attraction.
- For example, magnesium (Mg) forms stronger metallic bonds than potassium (K) because magnesium has smaller ions and more electrons in its sea.

iii. Arrangement of Metal Ions:

The closer packing of metal ions in the crystal lattice results in stronger metallic bonding. The closer the ions are to each other, the greater the electrostatic attraction between the metal ions and the delocalized electrons.
- Face-Centered Cubic (FCC) and Hexagonal Close-Packed (HCP) structures pack the ions very efficiently, contributing to stronger metallic bonds in metals like copper (Cu) and aluminum (Al).
- In contrast, Body-Centered Cubic (BCC) structures, which are less efficient in packing, tend to have slightly weaker metallic bonds.

iv. Factors Affecting the Strength of Metallic Bonds:

Factor	Effect on Bond Strength	Examples
Number of Delocalized Electrons	More electrons = stronger bonds.	Transition metals (e.g., Cu, Fe)
Size of Metal Ions	Smaller ions = stronger bonds.	Magnesium (Mg) > Potassium (K)
Arrangement of Metal Ions	Tighter packing = stronger bonds.	FCC and HCP structures (Cu, Al)

2.4.A.2 Interstitial Alloys: Smaller Atoms in Larger Atomic Spaces:

1. Definition of Interstitial Alloys: Interstitial alloys are a type of alloy where smaller atoms (often non-metals or elements with smaller atomic radii) occupy the interstitial spaces (or voids) between the metal atoms in the crystal lattice. This contrasts with other alloys, where atoms of similar or nearly similar radii replace metal atoms in the lattice or form solid solutions.

i. Formation of Interstitial Alloys:

Atomic Radii Difference: In interstitial alloys, the key characteristic is the difference in the sizes of the atoms involved. Typically, the atoms of the alloying element are smaller than the metal atoms in the host metal. These smaller atoms fit into the interstitial spaces (gaps or voids) between the larger metal atoms in the metal lattice.
How They Form:
- The host metal has a regular crystal structure, with metal ions arranged in a specific lattice (e.g., body-centered cubic, face-centered cubic, etc.).
- The small atoms (the alloying elements) fit into the spaces between the larger metal atoms. These small atoms don’t replace any of the metal atoms; instead, they occupy the gaps between the larger metal atoms.

2. Properties of Interstitial Alloys:

Alloy	Host Metal	Smaller Atom (Alloying Element)	Atomic Radii Comparison	How it Forms	Effect on Properties
Steel	Iron (Fe)	Carbon (C)	Iron has a larger atomic radius than carbon	Carbon atoms fit into the interstitial spaces between the iron atoms	Increases hardness and strength by disrupting the lattice structure.
Hydrogenated Palladium	Palladium (Pd)	Hydrogen (H)	Palladium has a larger atomic radius than hydrogen	Hydrogen atoms occupy the interstitial spaces between palladium atoms	Alters the properties of palladium, potentially improving certain characteristics like hydrogen absorption capacity.
Carbides and Nitrides	Tungsten (W)	Carbon (C)	Tungsten has a larger atomic radius than carbon	Carbon atoms fit into the interstitial spaces of the tungsten lattice	Enhances hardness and strength due to the disruption of the tungsten lattice.

3. Comparison with Substitutional Alloys:

Aspect	Interstitial Alloys	Substitutional Alloys

Atomic Size Difference

Large difference (smaller atom occupies interstitial space)

Smaller difference (atoms replace each other in the lattice)

Atomic Arrangement

Smaller atoms fit in voids between the host metal atoms

Alloy atoms substitute for metal atoms in the crystal lattice

Effect on Lattice

Distorts the regular lattice structure, increasing hardness

Can cause slight distortion or maintain regular structure

Effect on Properties

Increases hardness/strength, reduces ductility

Can improve strength, corrosion resistance, and ductility

Examples

Steel (Fe + C), Tungsten Carbide (W + C), Hydrogenated Palladium (Pd + H)

Brass (Cu + Zn), Bronze (Cu + Sn), Sterling Silver (Ag + Cu)

2.4.A.3 Substitutional Alloys: Atoms of Comparable Radius:

1. Formation and Structure:

i. Substitutional Alloy Formation:

- Occurs when atoms of two metals with comparable radii mix, forming a solid solution.
- The atoms of one metal replace those of the other in the lattice structure.

ii. Atomic Size Compatibility:

Atoms must have similar sizes (typically within 15% of each other) to fit properly into the crystal lattice without causing significant distortion.

iii. Example: Brass:

- Zinc atoms substitute for copper atoms in the brass lattice, forming a substitutional alloy.

iv. Resulting Properties:

- The substitution can alter properties like strength, hardness, and corrosion resistance, depending on the proportion and type of elements involved.

v. Effect on Lattice Structure:

- The alloy retains the basic crystal structure of the host metal, ensuring that the overall lattice remains intact but with altered properties due to the substitution.

vi. Applications:

- Common in materials like brass, bronze, and sterling silver, where different metals are combined to achieve specific mechanical or aesthetic properties.

2. Properties:

i. Impact on Strength:

- Substitutional alloys often exhibit improved mechanical properties, such as increased strength and hardness, due to the presence of different sized atoms disrupting the movement of dislocations in the lattice.

ii. Conductivity:

- The electrical conductivity can be affected by the type and amount of substitution. For example, in brass, the conductivity may be lower than pure copper due to the presence of zinc atoms.

iii. Appearance:

- Substitution of one metal for another can change the color or surface properties of the alloy. For example, brass (copper + zinc) has a yellowish color, while pure copper is reddish.

3. Example: Brass (Zinc Substituting for Copper):

Brass is a common example of a substitutional alloy, where zinc atoms replace copper atoms in the copper lattice.
Atomic Size: Zinc and copper have similar atomic radii, making them compatible for forming a substitutional alloy.
- Copper (Cu) has an atomic radius of about 1.28 Å.
- Zinc (Zn) has an atomic radius of about 1.34 Å.
This small difference in size allows zinc atoms to fit comfortably into the spaces in the copper lattice, replacing copper atoms.
i. Properties of Brass:
- The substitution of zinc for copper in brass alters its properties:
  - Strength: Brass is stronger and more durable than pure copper.
  - Color: Brass has a yellowish appearance, which makes it desirable for decorative uses.
  - Corrosion Resistance: Brass typically has improved resistance to corrosion compared to pure copper.

4. Comparison with Interstitial Alloys:

Property	Substitutional Alloys	Interstitial Alloys
Atomic Arrangement	Atoms of similar size replace each other in the lattice.	Smaller atoms occupy interstitial spaces between larger atoms.
Atomic Size Difference	Atoms have comparable radii (usually within 15%).	Atoms have significantly different sizes.
Examples	Brass (Zinc in Copper), Bronze (Tin in Copper)	Steel (Carbon in Iron), Hydrogen in Palladium
Effect on Lattice	Lattice maintains its overall structure but altered by substitution.	Lattice is distorted due to the presence of smaller atoms.
Impact on Properties	Alters strength, hardness, and conductivity.	Increases hardness, strength, and can affect conductivity.
Physical Appearance	Changes color, surface properties (e.g., brass is yellow).	May have little visible change; changes are due to atomic scale.
Common Applications	Decorative items, coins, electrical components.	Stronger materials like steel, hydrogen storage materials.

Metallic Solids (Old Content)

Metallic solid: are held together by the strong forces of attraction between the positively charged metal ions and delocalized electrons (electrons not associated with a single atom or molecule)
What accounts for differences in compound structures? The answer lies in the bonding

Comparing how Metals and Ionic Compounds Conduct Electricity

Metals → mobile valence electrons
Ionic compounds → mobile charged particles

Properties of Metals

Tendency to give up one or more electrons to form a positive ion → have low ionization energies
Malleable: easy to shape/bend rather than break bcuz adjacent layers of positive metal ions can move relative to one another while remaining in full contact with the electron sea
Ductile: pulled into wires
Shiny: reflect light bcuz free electrons can bounce back light at the same frequency
Solid at room temperature (except mercury which becomes liquid)
Metals conduct electricity and heat very efficiently because of the availability of highly mobile electrons
- Metal atoms in a metallic solid generate a “sea” of mobile valence that can easily move from one atom to the next and allow electricity to flow through; Electrons are excited from filled orbitals into the very near empty orbitals
  - If there is a large gap between the filled and the empty levels → electrons cannot be transferred easily to the empty conduction bands = poor conductor

Metal Alloys

Alloy: substance that contains a mixture of elements and has metallic properties → two types (difference is size of molecules)
Substitutional alloy: atoms of similar sizes
- Some of the host metal atoms are replaced by another element’s of similar size
Interstitial alloy: atoms of different sizes
- Formed when some of the holes in the metal’s structure are occupied by small atoms → presence of the interstitial atoms changes the properties of the host metal

Bond Polarity and Dipole Moments

Polarity: the difference in electronegativity (not the absolute value of electronegativity)
Dipole/having a dipole moment: molecule that has a center of positive charge and a center of negative charge (bcuz of unequal sharing of e-) → polar molecule
- Often has arrow pointing to the (-) charge center
Difference in electronegativity between atoms tells us how polar a compound is
- Atoms with high electronegativity with develop a partial negative charge
- Atoms with low electronegativity with develop a partial positive charge
Any diatomic molecule that has a polar bond will have a molecular dipole moment
With three or more atoms there are two considerations if will have a dipole moment
- There must be a polar bond & geometry cannot cancel it out
If any of the terminal elements are different → automatically a polar molecule (polar bonds are asymmetric)
- Unbalanced pull of electrons creates a partial (-) charge on one side of central atom and partial (+) charge on other → separation of charge forms a dipole = polar

Example: HF

Red indicates the most electron-rich area (the fluorine atom), and blue indicates the most electron-poor region (the hydrogen atom

Geometry and Dipole Moment

Some molecules have polar bonds but do not have a dipole moment (are nonpolar).
- - Occurs when bond dipoles are equal in magnitude and opposite in direction, canceling each other out and making the net dipole for the molecule zero
    - Ex: CO2 held by polar covalent bonds but is nonpolar
A molecule is nonpolar when all of the bond dipoles cancel each other out
- - Has equal pull of electrons and symmetrical geometry
  - Molecules composed of only nonpolar bonds are nonpolar, regardless of shape
Molecules in which the central atom is symmetrically surrounded by identical atoms are nonpolar, even if the bonds are polar
- - Equal attraction of electrons → net dipole is zero
Polar molecules may have some nonpolar bonds
Ex: of shapes that do cancel polarity of bonds
- - Ex: linear molecules with 2 identical bonds, planar molecules with 3 identical bonds 120 degrees part, tetrahedral molecules
Ex: of shapes that do NOT cancel polarity of bonds
- - Bent, trigonal pyramidal
→ Nonpolar: symmetrical and identical atoms
→ polar: unequal share of electrons

Energy Effects in Binary Ionic Compounds

Lattice energy: indicates how strong a bond is (but unlike bond energy is talking about ionic bonds)
- Determines how strongly the ions attract each other in the solid state
Lattice energy is stronger in ionic compounds with more highly charged ions and shorter ionic distance (smaller atoms) → the one with the stronger attraction

Example Question: Out of the following pairs of compounds, pick the one w/ greater lattice energy and explain

- (2)(-2) = -4 and (1)(-1) = -1 → MgO is the correct answer
- Explain: “MgO contains the more highly charged ion and thus the ions within that compound will experience a stronger attraction”
- Have same value for charged ions, but Na is smaller than Sr and O is smaller that I → Na2O is the correct answer
- Explain: “although these two compounds contain equally charged ions, sodium oxides have a shorter ionic distance thus the force of attraction is stronger”
Process: when given lattice energy question, write out the charges above each ion (ignore subscripts); then, if needed, compare sizes of atoms

Partial Ionic Character of Covalent Bonds

All bonds have a certain percent ionic character
- Compounds with more than 50% ionic character are usually ionic
Any compound that conducts an electric current when melted will be classified as ionic.

The Localized Electron Model

Localized electron (LE) model: assumes that bonds and molecules are formed from the overlapping of atomic orbitals on diff. atoms
Electron pairs are assumed to be localized on a particular atom or in the space between two atoms
- Lone pairs: belong to one atom/nuclei
- Bonding pairs: belong to 2 atoms/nuclei → found in the space between the atoms
  - Delocalization: electrons are free to move throughout the entire molecule; can be with either atom

Structure of Metals and Alloys Study Notes | AP Chemistry

Structure of Metals and Alloys Study Notes- AP Chemistry- Unit 2.4 -New Syllabus 2024-2025