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​Big Idea 2: Chemical and physical properties of materials can be 
explained by the structure and the arrangement of atoms, ions, or molecules and the forces between them. 

Video Resources
Bozeman Science
Practice Using the Ideal Gas Law Equation
​Dalton's Law of Partial Pressure
Molarity Calculations
Dilutions
​Very basic Lewis Structures (non-break octet)
​More basic Lewis Structres, mobile compliant
Essential knowledge 
All of the material below comes directly from the College Board and their publication, AP Chemistry, Course and Exam Description found on their website at the AP Chemistry course homepage.

2.A.1: The different properties of solids and liquids can be explained by differences in their structures, both at the particulate level and in their supramolecular structures.


2.A.2: The gaseous state can be effectively modeled with a mathematical equation relating various macroscopic properties. A gas has neither a definite volume nor a definite shape; because the effects of attractive forces are minimal, we usually assume that the particles move independently.

2.A.3: Solutions are homogenous mixtures in which the physical properties are dependent on the concentration of the solute and the strengths of all interactions among the particles of the solutes and solvent.

2.B.1: London dispersion forces are attractive forces present between all atoms and molecules. London dispersion forces are often the strongest net intermolecular force between large molecules.

2.B.2: Dipole forces result from the attraction among the positive ends and negative ends of polar molecules. Hydrogen bonding is a strong type of dipole-dipole force that exists when very electronegative atoms (N, O, and F) are involved.

2.B.3: Intermolecular forces play a key role in determining the properties of substances, including biological structures and interactions.

2.C.1: In covalent bonding, electrons are shared between the nuclei of two atoms to form a molecule or polyatomic ion. Electronegativity differences between the two atoms account for the distribution of the shared electrons and the polarity of the bond.

2.C.2: Ionic bonding results from the net attraction between oppositely charged ions, closely packed together in a crystal lattice.

2.C.3: Metallic bonding describes an array of positively charged metal cores surrounded by a sea of mobile valence electrons.

2.C.4: The localized electron bonding model describes and predicts molecular geometry using Lewis diagrams and the VSEPR model.

2.D.1: Ionic solids have high melting points, are brittle, and conduct electricity only when molten or in solution.

2.D.2: Metallic solids are good conductors of heat and electricity, have a wide range of melting points, and are shiny, malleable, ductile, and readily alloyed.

2.D.3: Covalent network solids generally have extremely high melting points, are hard, and are thermal insulators. Some conduct electricity.

2.D.4: Molecular solids with low molecular weight usually have low melting points and are not expected to conduct electricity as solids, in solution, or when molten.
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