AP Chemistry Unit 2 - Molecular and Ionic Compound Structures and Properties

University:
Kentucky State University
Course:
CHE 101 | General Chemistry I
Academic year:

2022
Views:

39

Pages:

3
Author:

Emilee Weeks

Unit 2 - Molecular and Ionic Compound Structures and Properties Coulomb’s law - Ions have stronger attractive force if they have larger charge and smaller ionic radius(smaller distance between ions) ● Bond energy - smaller bonds require more energy to break(double bonds have more bond energy than single bonds) ● Lattice energy - energy released when ions form a crystalline solid, smaller ions with higher charge have more lattice energy Ionic bonding - metal and nonmetal, conducts electricity as liquid/aqueous Covalent bonding - nonmetal and nonmetal, conducts electricity as liquid/aqueous Metallic bonding - sea of electrons move freely throughout a metal, conducts electricity as liquid/aqueous/solid Substitutional alloy - the atoms of both metals are roughly the same size Interstitial alloy - the atoms of one metal is smaller than the atoms of the other, so the smaller atoms fill interstitial space Bond Energies: To break bonds, energy must be added to the system and to form bonds, energy is released. The formula for this is Bond Energy also depends on the type of bond: Single Bond < Double Bond < Triple Bond The shorter the bond length, the greater the bond energy due to increased attraction between two atoms. Lewis Dot Structures ● Shows how valence electrons are arranged among atoms in a molecule. There are multiple rules for Lewis Structure: ● Duet Rule: This occurs between Hydrogen and Helium and it shares 2 electrons. ● Octet Rule: This occurs when each atom has 8 electrons. However, there are exceptions to this rule, which would be discussed later. ● Boron and Beryllium form an incomplete octet. Boron can only have 6 electrons, and Beryllium can only have 4 electrons. ● There are also elements that exceed the octet rule. Atoms that exceed row 3 or above, can form an expanded octet. For example, Sulfur, Phosphorous, Arsenic can all form an expanded octet. ● To determine formal charge, count each bond(double bond=2) and valence electron(lone pair=2). Subtract this number from the actual number of valence electrons on the periodic table. Try to get formal charge as close to 0 for a molecule and as close to the charge for an ion, ● Central atoms beyond period 3 can be hypervalent and violate the octet rule ● Resonance structures are when there is more than one way to represent a compound in a Lewis structure. Even though in a resonance structure there is more than one way to represent a compound, technically neither of them is right because in the actual molecule, the bond length is an average of both structures. NO3- is a good example with 3 resonance structures: VSEPR = Valence Shell Electron Pair Repulsion When drawing a Lewis structure, we only consider the valence electrons. Electrons in a bonding domain or in a nonbonding domain normally exist in pairs. Electron domains tend to repel each other in a way that is as far apart from each other in 3-D space as possible. An electron domain is a bond, but you don't count double or triple bonds as extra domains. Lone pairs on the central atom impact the shape. Sigma and pi bonding - each single bond is a sigma bond, any other bonds are pi bonds(triple bond has a sigma and two pi bonds)