AP Chemistry Unit 2:
Molecular and Ionic Compound
Structure and Properties

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2.1 Types of Chemical Bonds

• Ionic Bonding

• A bond type that transfers electrons to one another, usually a metal and a nonmetal.

• Atoms that lose or grain an electron are called cations (usually metals) or anions (usually nonmetals), respectively.

• Forms when an atom with low ionization energy (low energy required to remove an electron) reacts with an atom with high electron affinity (high energy amounts released when gaining an electron)

• Covalent Bonds

• A bond type that shares electrons, usually between two nonmetals.
• Polar Covalent Bond

• Unequal sharing of electrons, with one end positive and the other negative.

• Nonpolar Covalent Bond

• Equal sharing of electrons

• Electronegativity on Polarity

• Metallic Bonds

• This bond type is formed between metals with delocalized (easily movable) electrons because they are passing the electrons around. Metallic elements do not want electrons and are usually at the bottom left of the periodic table.
• Delocalized electrons are why metals are good conductors of both heat and electricity.
• Delocalized electrons are always valence electrons.
• In the image below, electrons are chaotic and are floating around between the metal ions.
  
  Source: byjus

2.2 Intramolecular Force and Potential Energy

• Bond Energy (Bond Order)

• Bond Order is also known as the # of bonds.

• Bond Length

• The distance where the potential energy between atoms is at its minimum.

• Intermolecular Distance Graph

• The lowest point (bond energy) is where the atoms are going to bond because it is where repulsion and attraction are balanced. The energy is negative because bonding releases energy (exothermic).
• Moving to the left of the bond energy causes the atoms to get too close together, pushing each other away (repulsion) and becoming unstable.
• Moving to the right of the bond energy causes the atoms to get too far apart, and the energy value is close to zero because one atom would be unable to attract the electrons around the other atom.
  
  Source: stackexchange

• Lattice Energy: The energy released when ions form as an Ionic Solid

• Very similar to Bond Energy, it is defined as the change in energy that takes place when gaseous ions are combined to form an ionic solid. Combining ions will release energy.
• The process by which a bond is formed is negative (exothermic) because electrons are usually attracted to lower energy levels. The process by which a bond is broken is positive (endothermic) because energy is required to break bonds.
• Because of the multiple ions, it is composed of multiple Coulomb Law interactions, which essentially represent the energy required to completely separate one mole of an ionic compound into its gaseous ions.
• 

• E = Lattice Energy
• k = Positive Constant
• Q1, Q2 = Charge of Ions
• r  = Distance between the center of ions (in nanometers)

• You will not need to calculate using the formula, but you will need to be able to determine between two substances if one substance has a higher or lower lattice energy than the other.
• Moving → or ↑ (Lattice Energy Increases)

• Lattice energy increases when the ion charge is increased with a decreased distance between them.

2.3 Structure of Ionic Solids

• Ionic Structure (Also called Crystal Lattice and Lattice Structure)

• The structure is held by strong electrostatic forces (attractive forces) because of the opposite charges of the cations and anions.
• They have very high melting and boiling points because they are aligned very tightly in a brittle structure, making them poor conductors.

• Ionic Structure Properties

• High melting and boiling points
• Poor conductors of electricity in solid form but are good conductors of electricity in liquid form. This is because ionic solids do not have free movement of ions.
• Hard and brittle

• Ionic Bonds

• The relative strength of different ionic bonds can be estimated using Coulomb's Law. The force is proportional to each of the charges and inversely proportional to the distance squared.
• Increasing the distance between changes reduces the strength of the Coulombic attraction.

2.4 Structure of Metals and Alloys

• Solid Properties

• Alloys

• Alloys possess metallic properties and are composed of two or more elements that can be either homogeneous or heterogeneous.
  
  Source: libretexts.org
• Substitutional Alloy

• The similar-sized atoms will replace each other throughout the structure.

• Interstitial Alloy

• The smaller atoms fit into the gaps between the larger atoms.

2.5 Lewis Diagrams

• Lewis Diagrams (Also known as Lewis Structures)

• A type of localized electron model that is used to show how valence electrons and bonds are arranged.
• Octet Rule

1. Lewis Structures are based on the Octet Rule, which most atoms follow: Atoms combine so that they each have 8 electrons in their valence shells, allowing them to be stable.
2. Exceptions:

• Odd Number of Electrons

• Example: NO (11 electrons)

• Atoms with more than 8 electrons

• 3rd Row and heavier elements

• Atoms with less than 8 electrons

• Hydrogen (H), Helium (He), Beryllium (Be), and Lithium (Li) because they contain 1, 2, 2, and 3 electrons, respectively.
• Boron (B) and Aluminum (Al) can be full with 6 electrons

• Lone Pairs: Unshared (lone) pair (two) of electrons on a single atom
• Bonding Pair: Shared pair (two) of electrons
• Steps to Drawing a Lewis Structure

1. Identify Central Atom
2. Find the total number of valence electrons (remember charge)
3. Make a framework for the molecule
4. Make bonds between the atoms
5. Add electrons to the atoms until they fulfill the Octet Rule
    Tip: If the central atom does not have 4 electron pairs, try double or even triple bonding

2.6 Resonance and Formal Charge

• Resonance

• This occurs when a Lewis Structure can be drawn in multiple valid ways with the bindings between the same elements. In the example, the structure requires an oxygen molecule to have a double bond, but it doesn't matter which oxygen molecule has it, making it a resonance structure.
• On the AP Exam, you will need to draw all the resonance structures with a double-sided arrow between each structure.
• Bond Order

• Single bonds have a bond order of 1, and a double bond has a bond order of 2.
• Formula: Total Bond Order (single bonds + double bonds)/ # of resonance forms

• Example:

• Bond Order:  or 1.33

Source: Toppr

• Formal Charge

• Used to predict the correct placement of electrons across different elements. Essentially the same thing as a resonance structure, but for different elements
• Formal charges will add up to the polyatomic ion's charge
• Calculating formal charges for each atom:

• Valence electrons - ( lone pairs + bonds)

• The fewer number of atoms with a formal charge of zero, the more likely the structure will be because molecules try to achieve formal charges as close to zero as possible.
• Same number of zeros:

• The negatively charged atom with the highest electronegativity would be most likely to form.
• Using the example, we can determine that A and B both have two formal charges of zero, making them the most likely structures that would be formed. C only has one formal charge of zero, making it the most inaccurate placement of electrons. We know that Nitrogen (N) is more electronegative than Sulfur (S), making B the most accurate placement of electrons.

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• Example: SCN−

2.7 VSEPR and Hybridization

• VSEPR Diagram

• VSPER, standing for Valence Shell Electrons Pair Repulsion, is used to predict molecular structure by minimizing electron pair repulsions
• Steps to Predict

• VSEPR Structures:

• You should memorize this chart
  (excluding the hybridization for electron groups 5 and 6)
  
  Source: mst.edu

• Predict the VSEPR Model

1. Draw Lewis Structure
2. Space out the electron pairs
3. Using the way electron pairs are shared, determine the position of atoms
4. Name the molecular structure (Linear, Bent, etc.)

• Lone pairs take up more space than bonding pairs, causing angles between bonding pairs to compress.
• Double and Triple bonds are treated the same way as single bonds when predicting VSEPR structures. However, they do have more repulsive strength and take up slightly more space than single bonds,

• Hybridization

• When atomic orbitals overlap, forming special (hybridized) orbitals for bonding.
• The electron orbitals overlap each other to fill subshells and go to a lower energy state. Atoms want to give the molecule the minimum amount of energy.
• The number of hybridized orbitals is the total number of orbitals used to create it. sp3 (1s orbital & 3p orbitals) = 4 hybridized orbitals
  
  Source: matthewrkennedy

• Bonding

• Sigma (σ) Bonds

1. Single bonds are Sigma bonds
2. Sigma bonds are localized, meaning they do not move around

• Pi (π) Bonds

1. The Second and Third bonds are Pi bonds and can be above or below a Sigma bond.
2. Resonating (Lewis Structures that can be drawn in multiple valid ways with the bindings between the same elements) Pi bonds are delocalized, meaning they can move around.
3. In triple bonds, the Pi bonds are on both sides of the Sigma bonds and are less strong.
4. More Pi bonds mean higher bond energy but shorter bond lengths.