AP chem review

Definition of Matter: Anything that has mass and occupies space.
Atoms: The smallest unit of matter retaining properties of an element; molecules are bonded atoms.
Moles and Molar Mass:
- Mole: The unit for the amount of substance; equals Avogadro’s number, 6.022 imes 10^{23} particles.
- Molar Mass: Mass of one mole of a substance in grams/mole, used to convert grams to moles:
- ext{# of moles} = rac{ ext{weight in grams}}{ ext{molar mass}}
- ext{# of molecules} = ext{moles} imes (6.022 imes 10^{23})

Function: Measures relative abundance of isotopes by separating ions by charge and weight.
- Process involves charging samples, using magnetic fields for separation, and measuring relative abundance.
- Results plotted with mass-to-charge ratio (m/z) and relative abundance.

Pure Substances: Consist of a single type of atom (element) or molecule (compound).
- Law of Definite Proportions: Fixed elemental ratio in compounds.
Mixtures: Composed of two or more substances; can vary in composition.
- Homogeneous Mixtures: Uniform composition.
- Heterogeneous Mixtures: Non-uniform composition.

Subatomic Particles:
- Protons and Neutrons: Mass of about 1.67 imes 10^{-27} kg.
- Electrons: Negligible mass; orbit outside the nucleus.
Atomic Number and Mass Number:
- Mass Number (A): Total of protons and neutrons.
- Atomic Number (Z): Number of protons, defining the element. Isotopes have the same Z but different A.
Electron Configuration:
- Electrons are arranged in shells and subshells; filled according to the Aufbau principle and Hund’s Rule.
- Electron count is determined by subshell distribution. Example:
- Carbon: 1s^2 2s^2 2p^2
Ions:
- Loss of electrons gives positive cations; gain of electrons gives negative anions.

Coulomb's Law: F = k rac{Q_1Q_2}{r^2}; impacts ionization energy based on distance and charge interaction.

Usage: Determines electron arrangement in atoms; involves measuring energy and relative abundance of electron emission.

Trends in the Periodic Table:
- Atomic radius, ionization energy, and electronegativity trends explained with electron configurations.
Periodic Families:
- Alkali Metals: Highly reactive, +1 charge.
- Alkaline Earth Metals: +2 charge, less reactive than alkali metals.
- Transition Metals: Variable charges, conductive and colorful.
- Chalcogens: -2 charge.
- Halogens: -1 charge, highly reactive.
- Noble Gases: Unreactive with full valence shells.

Ionic Bonds: Formed by electron transfer, stability achieved through full valence shells.
Duet and Octet Rules:
- First five elements stable with 2 electrons; others stable with 8.

Calculate mass ratios, balance chemical equations, or identify isoelectronic configurations as per questions derived from test structure.

Covalent Bonds: Electron sharing for stability, form individual molecules (examples: CO2).
Ionic Bonds: Electrostatic attractions between charged ions, form crystalline structures (Example: NaCl).
Metallic Bonds: Positive nuclei surrounded by a sea of electrons, account for conductive and malleable properties.

Electronegativity Scale: High values correspond to ionic nature (>1.7); low values correspond to covalent nature (<0.5).

Potential energy diagrams that illustrate bond lengths and energies important for covalent bonding concepts.

Lattice Structures: Stability due to optimization in charge interactions; yielding distinct properties for ionic compounds.

Drawing Lewis Structures: Follow the octet and duet rules; represent bonds and lone pairs effectively.
VSEPR Theory: Predict molecular shapes based on electron pair repulsion.

Access additional readings on atomic structures, interactions, bonding theories, and kinetic molecular theories for deeper understanding.

Intermolecular Forces (IMF): Attraction or repulsion between molecules, non-chemical bond forces.
- Types: London dispersion forces, dipole-dipole interactions, hydrogen bonds, affecting physical states and properties.

Distinction between solids (fixed structure), liquids (adapt to container shape), and gases (occupy any volume).

Introduced relationships among pressure (P), volume (V), and temperature (T) in regards to kinetic behaviors. Law formulas such as: PV = nRT.

Types of Changes: Define physical versus chemical changes.
Net Ionic Equations: Outline reactions whilst disregarding spectator ions.
Stoichiometry: Calculate molar quantities and relationships in chemical reactions.

Distinguish endothermic and exothermic reactions concerning energy shifts in reactions; calculate enthalpy changes following Hess's Law.

reactions are reversable
at equilibrium, the rate of the forward reaction is equal to the rate of the reverse reaction. This does not mean they are equal
Keq is constant at a given temp

LeChatelier’sPrinciple