1. Gases
1.1. Ideal gas law: PV = nRT
1.1.1. R = 0.08206 L*atm/mol*K
1.1.2. Follows from Simple Gas Laws:
1.1.2.1. Boyle's Law: P1V1 = P2V2
1.1.2.2. Charles's Law: V1/T1 = V2/T2
1.1.2.3. Avogadro's Law: V1/n1 = V2/n2
1.1.3. Molar volume of ideal gas at STP (standard temperature and pressure) = 22.4 L
1.1.4. density = PM/RT, where M = molar mass
1.1.5. to find molar mass, use n = PV/RT form of ideal gas law, then divide given g/mol
1.2. Kinetic molecular theory
1.2.1. Basic assumptions:
1.2.1.1. 1. The size of a particle is negligibly small. (assumes that the particles themselves have no volume)
1.2.1.2. 2. The average kinetic energy of a particle is proportional to the temperature in kelvins. (the higher the temperature, the faster everything moves)
1.2.1.3. 3. The collision of one particle with another (or the walls of the container) is completely elastic (energy is exchanged but not lost).
1.3. Pressure = Force/Area (collisions of gas particles on th surface of container)
1.3.1. 1 atm = 760 torr = 101,325 Pa(N/m^2) = 29.92 inHg = 1.013 bar
1.3.2. Partial pressures: each gas component in a mixture of gases acts indepently
1.3.2.1. Ptotal (total pressure) = Pa + Pb + Pc + ... (Dalton's Law)
1.3.3. Mole fractions (Xa)
1.3.3.1. Xa = na/ntotal
1.3.3.2. Pa (Partial pressure for gas 'a') = xa/Ptotal
2. Electron configuration
2.1. Valence electrons: those in the outermost principal energy level
2.1.1. last digit of group number
2.1.2. Important for bonding because they're held most loosely (atoms want to attain noble gas configuration)
2.1.2.1. those towards right side of table tend to GAIN (nonmetals, anions)
2.1.2.2. those to left tend to LOSE (metals, cations)
2.1.2.3. those in the middle - especially metalloids and nonmetals - tend to SHARE (covalent/polar covalent)
2.1.3. Periodic trends
2.1.3.1. Atomic radius = set of average bonding radii
2.1.3.1.1. nonmetals: 1/2 distance between 2 bonded atoms
2.1.3.1.2. metals, 1/2 distance between 2 adjacent atoms in a crystal lattice
2.1.3.2. Ionic radius = atomic radius of an ion
2.1.3.2.1. anions: LARGER THAN neutral
2.1.3.2.2. cations: SMALLER THAN neutral
2.1.3.3. Ionization energy = energy required to remove one electron from an atom/ion in gaseous state
2.1.3.4. Electron affinity = (delta)E associated with gaining an electron in the gaseous state, usually negative and an exothermic process
2.1.3.5. Effective nuclear charge = #protons - #core electrons
2.2. Ionization energy
2.3. Electron affinity (attract any electron)
2.3.1. ELECTRONEGATIVITY (attract electron within covalent bond)
2.3.1.1. types of compounds
2.3.1.1.1. Covalent: sharing electrons between nonmetals (which have high ionization energy). 0-0.4 ΔEN
2.3.1.1.2. Polar covalent: electrons shared unequally. 0.4-2.0 ΔEN
2.3.1.1.3. Ionic: metal becomes cation, nonmetal becomes anion. 2.0+ ΔEN
2.3.1.2. Structure of molecules
2.3.1.2.1. Lewis dot structures
2.3.1.3. Polarity
2.3.1.3.1. generally increases up and across periodic table
2.3.1.3.2. EN Pauling Scale: F = 4.0, most electronegative element.
2.3.1.3.3. Dipole moment: measurement of net molecular polarity
2.3.1.4. Composition, formulae
2.3.1.4.1. Stoichiometry
2.3.1.4.2. Empirical formula = relative # of atoms of each element
2.3.1.4.3. Molecular formula = actual # of atoms of each element
2.3.1.4.4. Structural = shows how atoms are connected/bonded
2.3.1.4.5. Bond energy = energy required to break the bond between 1 mole of atoms in the gas phase
2.3.1.4.6. Bond length = measured as the distance between bonded nuclei
2.4. Shielding: repulsion by electrons in lower-energy orbitals that screen valence electrons from the full effects of nuclear charge
2.5. Penetration: outer electron penetrates into the region of the inner ones and experiences greater nuclear charge and lower energy
2.5.1. Due to Coulomb's law: potential energy of two charged particles depends both on their charges and their separation
3. Composition of atoms
3.1. Mole concept
3.1.1. Avagadro's number
3.1.1.1. 6.022*10^23
3.1.2. Molar mass = g/mol = atomic mass in amu
3.2. Classifying matter
3.2.1. Pure substance: one particle, proportions do not vary between samples
3.2.2. Mixture: Proportions between samples vary
3.3. Atomic theory
3.3.1. Law of Conservation of Mass: In a chemical reaction, matter is neither created nor destroyed
3.3.2. Law of definite proportions: All samples of one compound have the same proportions of the constituent elements
3.3.3. Dalton's theory
3.3.3.1. Elements are composed of atoms
3.3.3.2. All atoms of one element have the same mass and unique properties
3.3.3.3. Atoms combine in simple, whole number ratios to form compounds (law of multiple proportions)
3.3.3.4. Atoms of one element cannot change into atoms of another element
3.4. Subatomic particles
3.4.1. amu = 1/12 mass of a 12-6 Carbon atom
3.4.2. isotope = an atom of the same element (# of protons) but a different # of neutrons
3.4.3. mass number = p + n
3.4.4. atomic mass: average mass of isotopes, weighted according to natural abundance - sum of (fraction of isotope n)*(mass of isotope n)
4. Bohr model
4.1. Rydberg equation ΔE = -2.18*10^-18 J(1/nf^2 - 1/ni^2)
4.2. Quantization
4.2.1. Electron as standing wave
4.2.1.1. Atomic orbitals
4.2.1.1.1. Orbital: probability distribution map, describes likely position of electron
4.2.1.1.2. Aufbau principle
4.2.1.1.3. Hund's Rule
4.3. Quantum numbers
4.3.1. n (energy level)
4.3.1.1. "Principal" number, determines overall size and energy
4.3.2. l (subshell)
4.3.2.1. "Angular momentum", integers up to n-1, determines shape
4.3.3. m_l (orbital)
4.3.3.1. "Magnetic", integers =-l, steps of 1, ... +l, specifies orientation
4.3.4. m_s (electron spin)
4.3.4.1. "Spin", +1/2 or -1/2
4.3.4.2. PAULI EXCLUSION PRINCIPLE: no two electrons in an atom can have the same four quantum numbers
4.4. Emission spectrum
4.4.1. Electron jumps to an excited state, then relaxes back to ground state, releasing the same energy in brightly colored lines
4.5. Absorption spectrum
4.5.1. Dark lines show which wavelengths (and corresponding energies) are missing due to absorption
5. Energy
5.1. = the capacity to do work (a force acting through a distance)
5.1.1. kinetic: associated with motion
5.1.2. potential: associated with position or composition
5.2. Law of conservation of energy: it's neither created nor destroyed
5.3. Electromagnetic radiation
5.3.1. Light
5.3.1.1. Photoelectric effect (many metals emit electrons when light shines on them)
5.3.1.1.1. Threshold frequency
5.3.2. Oscillating waves
5.3.2.1. Amplitude
5.3.2.2. Wavelength (λ)
5.3.2.3. Frequency (v)
5.3.2.3.1. # of cycles that pass through a stationary point per sec
5.3.2.3.2. = c/λ
5.3.3. E = hv (hc/λ)
5.3.3.1. h = Planck's constant
5.3.3.2. c = speed of light
5.3.3.3. λ = wavelength
5.4. Thermodynamics
5.4.1. Thermochemistry
5.4.1.1. Calorimetry
5.4.1.1.1. Bomb calorimeter - constant volume, w = 0
5.4.1.1.2. Coffee cup calorimeter - constant pressure, measures ΔH
5.4.1.2. System vs. Surroundings
5.4.1.2.1. System: whatever is under investigation, usually a chemical reaction
5.4.1.2.2. Surroundings: the rest of the universe
5.4.1.2.3. Internal energy change:
5.4.1.3. Thermal equilibrium: the point at which there is no more net transfer of heat between a system and its surroundings
5.4.2. Enthalpy ΔΗ: amount of heat absorbed/evolved in the reaction at constant pressure
5.4.2.1. Endothermic and exothermic reactions
5.4.2.1.1. +ΔH signifies endothermic reaction for the system
5.4.2.1.2. -ΔH signifies exothermic reaction for the system
5.4.2.2. Heat: flow of energy caused by a temperature difference
5.4.2.2.1. Heat capacity
5.4.2.2.2. Calculated with q = mCsΔT (can be read as "em-cat")
5.4.2.3. Temperature
5.4.2.4. ΔH = ΔΕ + ΡΔV = qp, heat at constant pressure
5.4.2.4.1. Calculations of ΔHrxn
5.4.3. State functions: value depends only on state of system, not on now it got there
5.4.3.1. Energy, pressure, volume are all state functions