unit 3 ap chem

Bonds are based on attractions and repulsions between charged particles.
Coulombs Law governs the strength of bonding:
- F = (\frac{q_1 q_2}{r^2})
Full charges on ions (e.g., Na+, Cl-) provide strong interactions, while partial charges ((\delta^+, \delta^-)) lead to weaker interactions.

Van der Waals Forces: Include London Dispersion Forces (LDF's), dipole-dipole interactions, and hydrogen bonding.
London Dispersion Forces: Weak, temporary dipoles created by shifting electron clouds; found in all substances but significant in larger atoms with greater electron cloud fluctuation.
Permanent Dipoles: Occur when there’s a significant electronegativity difference between atoms in a molecule. Typically stronger than LDF's.
- Example: Acetone possesses permanent dipoles, increasing boiling point compared to similar molecular weight but non-polar substances.
Hydrogen Bonds: Strong interactions formed when hydrogen is bonded to electronegative atoms (N, O, F).

Hydrogen bonding results in higher boiling points compared to substances relying solely on LDF's.
Ionic Compounds: High melting/boiling points; brittle due to lattice structures.
Covalent Network Solids: High melting points due to strong covalent bonds; typically non-conductive as electrons are fixed in place.
Metallic Bonds: Delocalized electrons allow for conductivity and malleability.

Solids are most stable at standard temperature and pressure (25°C, 1 atm).
- Crystalline Solids: Orderly structure with defined melting points.
- Amorphous Solids: Lack a defined structure.

Strong interactions at short distances; the melting points influenced by ion size and charge.
Conductivity: Ionic solids do not conduct electricity in solid states; they must be melted or dissolved in a solvent.

Mainly held together by LDF's; solubility is affected by molecular size and intermolecular forces.
- E.g., I2 is solid due to strong LDF's, while CO2 is gaseous due to weak LDF's.

Molecules are closely packed but have more freedom than solids; they can move and collide.
Liquids have a constant volume but take the shape of their container.
Example: Water expands when freezing, forming a hexagonal structure due to hydrogen bonds.

Gases have no definite shape or volume, are compressible, and fill their container.
The Ideal Gas Law and Kinetic Molecular Theory explain gas behaviors under various pressures and temperatures.

Boyle’s Law: (PV = \text{constant}) (inversely proportional)
Charles’s Law: (V/T = \text{constant}) (directly proportional)
Avogadro’s Law: (V/n = \text{constant}) (directly proportional)
Ideal Gas Equation: (PV = nRT)
- R = 0.08206 L atm mol^-1 K^-1 at standard conditions.

Gases consist of many molecules in random motion.
The volume of gas molecules is negligible compared to the volume of their container.
No significant attractive/repulsive forces exist between gas molecules.
Energy is transferred during collisions, but average kinetic energy remains constant.
The average kinetic energy varies with temperature.

Real gases deviate from ideal behavior due to the finite volume and intermolecular attractions.
Deviations are significant under high pressure and low temperature where attractive forces dominate.

Solution: Homogeneous mixture of solute and solvent; solute particles disperse uniformly through a solvent.
Miscibility: The ability of two liquids to mix; governed by their intermolecular interactions.

Solute-solvent interactions vs. solute-solute and solvent-solvent interactions.
Polar solvents can dissolve ionic compounds; non-polar solvents dissolve nonpolars.

Absorption and emission of light provide insights into atomic structure and interactions.
Different energy regions (visible, infrared, ultraviolet) correspond to different molecular transitions.

Absorbance is proportional to concentration and path length (A = εbc).
Calibration plots help determine unknown concentrations based on absorbance measurements.