BV Module1

• States of matter:

  • Solid: definite shape, non-compressible
  • Liquid: indefinite shape, assumes container shape, non-compressible
  • Gas: indefinite shape, compressible
• Plasma/Ionized Gas:

  • 4th state of matter, most abundant
  • Contains protons and electrons
  • Examples: ionized Ne light, Aurora, Stars, Sun
• Phase Change:

  • Melting (Solid to Liquid)
  • Freezing (Liquid to Solid)
  • Evaporation (Liquid to Gas)
  • Condensation (Gas to Liquid)
  • Sublimation (Solid to Gas)
  • Deposition (Gas to Solid)
  • Recombination (Plasma to Gas)
  • Ionization (Gas to Plasma)
• Matter Classification:

  • Pure substance:
  • Element: simplest form of substance
  • Compound: 2 or more chemicals united
  • Mixture:
  • Homogeneous: 1 phase, solution
  • Heterogeneous: 2 phases, suspension, colloid
• Intrinsic and Extrinsic Properties:

  • Extrinsic Property (Dependent): length, mass/weight, volume, pressure, entropy, enthalpy, electrical resistance
  • Intrinsic Property (Independent): density/SpGr, viscosity, velocity, temperature, color
• Fundamental Chemistry Laws:

  • Law of Conservation of Mass/Matter: Mass/Matter is always constant
  • Law of Definite/Constant Proportions: Chemical compounds contain exact proportions of elements
  • Law of Multiple Proportion: Elements form compounds in fixed whole numbers
  • Law of Combining Weights: Proportions by weight in chemical reactions
• Atomic Structure:

  • Democritus: Atomos, "Indivisible"
  • John Dalton: Matter is made up of atoms, elements composed of indivisible atoms
  • J.J. Thompson: Plum Pudding model, electrons in a positive framework
  • Ernest Rutherford: Nuclear model, atom is mostly empty with positive particles in the nucleus
  • Neil Bohr: Planetary model
  • Erwin Schrodinger: Quantum/Mechanical model, estimates probability of finding an electron in a certain position
• Chemical Bonds:

  • Molecule: aggregate of 2 or more atoms held together by chemical bonds
  • Ions: with net positive or negative charge
  • Empirical formula: simplest whole number ratio
  • Forces of Attraction: Intermolecular forces, H-bonding, Dipole, Induced Dipole
  • Intramolecular forces: Covalent, Ionic
  • Covalent Bonding: sharing of electrons, lone pair
  • Ionic Bonding: transfer of electrons
  • Valence Shell Electron Pair Repulsion (VSEPR) theory: predicts molecular geometry
• Reaction Types:

  • Synthesis/Combination/Direct Union: A + B → AB
  • Decomposition/Analysis
• Bonding:

  • Molecular orbitals can be either bonding (lower energy and stable) or antibonding (higher energy and unstable)
• Reaction types:

  • Synthesis/Combination/Direct Union: A + B → AB
  • Decomposition/Analysis: AB → A + B
  • Complete combustion example: CH4 + O2 → CO2 + H2O
  • Incomplete combustion example: CH4 + O2 → CO + C (5) + H2O
  • Single Displacement: AB + X → AX + B
  • Double Displacement/Metathesis/Exchange: AB + CD → AC + BD
  • Neutralization example: NaOH + HCl → NaCl + H2O
  • Precipitation example: AgNO3 + NaCl → AgCl2 + NaNO3
  • AgCl2 is a white curdy precipitate
• Reactivity series:

  • Metals: Li > K > Ba > Ca > Na > Mg > Al > Mn > Zn > Cr > Fe > Cd > Co > Ni > Sn > Pb > H2 > Cu > Ag > Hg > Pt > Au
  • Nonmetals based on electronegativity: F > Cl > Br > I
  • Examples of reactions: Co + MgCl2 → NR, Zn + CuSO4 → ZnSO4 + Cu, NaBr + Cl2 → NaCl + Br
• Nomenclature of inorganic compounds:

  • Covalent compounds examples: CO (Carbon monoxide), SiO2 (Silicon dioxide), N2O (Dinitrogen monoxide), CCl4 (Carbon tetrachloride)
  • Ionic compounds example: Pb(NO3)4
  • Classical name for Pb(NO3)4: Plumbic nitrate
  • Stock name for Pb(NO3)4: Lead(IV) nitrate
• Monoatomic ions:

  • Monovalent ions: +1 = Group 1 (H, Li, Na, K, Ag), +2 = Group 2 (Be, Mg, Ca, Sr, Ba, Zn, Cd), -2 = Group 6A (Oxide, Sulfide), -1 = Group 7A (Fluoride, Chloride, Bromide, Iodide)
  • Multivalent ions: +1, +2 = Hg, Cu, +1, +3 = Au, +2, +3 = Fe, Co, Ni, +3, +5 = Bi, Sb
• Polyatomic ions:

  • Oxyanions examples:
  • ClO-: Hypochlorite, Hypochlorous acid (HClO)
  • ClO2-: Chlorite, Chlorous acid (HClO2)
  • ClO3-: Chlorate, Chloric acid (HClO3)
  • ClO4-: Perchlorate, Perchloric acid (HClO4)
  • NO2-: Nitrite, Nitrous acid (HNO2)
  • NO3-: Nitrate, Nitric acid (HNO3)
  • SO32-: Sulfite, Sulfurous acid (H2SO3)
  • SO42-: Sulfate, Sulfuric acid (H2SO4)
  • PO43-: Phosphate, Phosphoric acid (H3PO4)
  • Naming conventions: ate for common form, ite for -1 O to ate form, hypo...ite for -1 O to ite form, per...ate for +1 O to ate form
• Mole relationships:

  • Avogadro’s number: 1 mole = 6.022 x 10^23 atoms/molecules
  • Example calculation of number of NaOH atoms using Avogadro’s number
• Gas laws:

  • Boyle's/Mariotte law: P1V1 = P2V2 or P ∝ 1/V
  • Charles' law: T1V1 = T2V2 or V ∝ T
  • Gay-Lussac's law: P1T1 = P2T2 or P ∝ T
  • Combined gas law: P1V1T1 = P2V2T2
  • Ideal gas law: PV = nRT
  • Avogadro’s principle: Equal volumes of different gases have the same number of moles at STP
  • Dalton’s law of partial pressures: Total pressure in a mixture is the sum of the partial pressures of each gas
  • Graham’s law: Rate of effusion and speed of gas are inversely proportional to the square root of their density
• • Fick’s 1st law: Diffusion rate of liquid or gas is directly proportional to the concentration gradient
  • Henry’s law of gas solubility: Pressure is proportional to solubility
• Temperature conversions:

  • °C to °F: (°F - 32) / 1.8
  • °F to °C: (°C x 1.8) + 32
  • K to °C: °C + 273.15
• Solution:

  • Solute + Solvent
  • Colligative properties: Vapor pressure lowering, Boiling point elevation, Freezing point depression, Osmotic pressure
• Thermodynamics:

  • Study of energy conversion/transformation in the universe
  • Parts of the universe: Open System, Closed System, Isolated System
  • Quantum numbers: Principal Quantum Number, Azimuthal/Angular Momentum, Magnetic Quantum Number, Magnetic Spin
  • Quantum theories: Pauli’s exclusion theory, Heisenberg’s uncertainty theory, Hund’s rule
• Gas laws:

  • Boyle's/Mariotte law, Charles' law, Gay-Lussac's law, Combined gas law, Ideal gas law, Avogadro’s principle, Dalton’s law of partial pressures, Graham’s law, Fick’s 1st law, Henry’s law of gas solubility
• Temperature conversions:

  • °C to °F, °F to °C, K to °C
• Thermodynamics:

  • Study of energy conversion/transformation in the universe
  • Parts of the universe: Open System, Closed System, Isolated System
• Types of Systems:

  • Open System: allows exchange of energy and matter
  • Closed System: allows exchange of energy but not matter
  • Isolated System "Adiabatic Walls": does not allow exchange of both energy and matter
• Surroundings: everything outside the system
• Path Dependence:

  • State Function: independent, depends only on initial & final states of the system
  • Examples: Enthalpy (H), Internal energy (U), Gibb’s Free Energy (G), Entropy (S)
  • Non-State Function: dependent, includes Work and Heat
• Zeroth Law:

  • If two systems are in thermal equilibrium with a third system, they must be in thermal equilibrium with each other
• Laws of Thermodynamics:

  1. 1st Law: Law of conservation of Energy

  • Energy is neither created nor destroyed but can be transformed from one form to another
  • Enthalpy (H) = U, P, V
  • Hess' Law: ∆ H is independent of reaction/steps that occurred
  • q = Heat
  • ∆ H = (+) → heat is absorbed; COLD (endothermic)
  • ∆ H = (-) → heat is released; HOT (exothermic)

  2. 2nd Law: Law of Entropy

  • No way but UP
  • Total entropy can never decrease over time for an isolated system
  • Entropy (∆ S) = measure of system’s thermal energy per unit temperature; degree of disorderliness or randomness
• • ∆ S = (+) → spontaneous; increase (irreversible)
  • ∆ S = (-) → non-spontaneous; constant (reversible)
  • ∆ H does not predict spontaneity

  3. 3rd Law:

  • Entropy of a perfect, solid, crystalline substance is zero at absolute 0 temperature
  • Gibb’s free energy (∆ G):
  • Combines enthalpy and entropy
  • ∆ G = ∆ H - T ∆ S
  • ∆ G < 0 (-) → spontaneous
  • ∆ G > 0 (+) → non-spontaneous
  • ∆ G = 0 → equilibrium (no more work to be done)
• Chemical Kinetics:

  • Study of reaction rates and reaction mechanism
  • Reaction Rate (M/s): Change in concentration of a reactant or product concentration with time
  • Rate Law: Expresses relationship of the rate of reaction to the rate constant (K) and concentration of reactants raised to some power
  • Rate = K [A] x [B] y (xy = order of reaction)
  • Factors Affecting Reaction Rate:
  • Nature of Reactants: ↑ reactivity ↑ reaction rate (faster)
  • Concentration of Reactants (except Zero order): ↑ concentration ↑ reaction rate
  • Catalyst (Enzyme - Michaelis Menten Kinetics): ↑ reaction rate
• • Surface Area: ↑ SA ↓ particle size ↑ reaction rate ↓ reaction time
  • Temperature: ↑ Temp ↑ KE ↑ mobility of molecules ↑ collision ↑ reaction rate
• Chemical Equilibrium:

  • Law of Mass Action: reaction rate proportional to the product of the concentrate of the reactants to the power of its coefficient in a balanced equation
  • Le Chatelier’s Principle: If an external stress is applied to a system at equilibrium, the system adjusts to partially offset the stress as it reaches a new equilibrium
  • Acids and Bases:
  • Acids: Taste Sour, pH < 7, + Litmus paper turns Red
  • Bases: Taste Bitter, pH > 7, + Litmus paper turns Blue
  • Acid-Base Formula:
  • For Weak Acids & Bases (with constant):
  • pH = - log[H +], pKa = - log[ka]
• • pOH = - log[OH -], pKb = - log[kb]
  • pH + pOH = 14, pKa + pKb = 14
  • Common Ion Effect: Addition of compound having an ion in common with the dissolved substance will result in equilibrium shift, suppressed ionization of the dissolved substance, and pH change
  • Henderson-Hasselbalch/Buffer pair equation: For buffer solutions (WA + C)
• Buffer pair equation:

  • HAc + Ac-
  • NH3 + NH4+
  • Weak acids: pH = pKa + log salt acid
  • Weak bases: pH = pKb + log base salt
  • pH = pKa (@ half neutralization point)
  • Buffer solution has the ability to resist changes in pH upon addition of small amounts of either acid or base
  • Weak acid and its conjugate base (salt of weak acid)
  • Weak base and its conjugate acid (salt of weak base)
• Solubility product constant (Ksp):

  • ↑ Ksp = ↑ solubility
  • Solubility (g/L): Number of grams of solute dissolved in 1L of saturated solution
  • Molar solubility (mol/L): Number of moles of solute dissolved in 1L of saturated solution
  • Predicting formation of precipitate formation (Q ion product constant):
  • Q < Ksp → unsaturated
  • Q = Ksp → saturated
  • Q > Ksp → supersaturated
• Electrochemistry:

  • Study of the production of electricity from energy released during spontaneous and nonspontaneous chemical reactions
  • Spontaneous: Voltaic cells/galvanic cells, REDOX reaction (Anode - Oxidation; Cathode - Reduction), Electrons migrate from Anode → Cathode
  • Nonspontaneous: Electrolytic cells, Electric current is applied to remove e- and transfer to another cell (Electroplating)
• Periodic Table:

  • Antoine Lavoisier: First extensive list of elements (~33), Metals vs Nonmetals
  • Metal properties: Oxides are basic, good reducing agents, conductors, malleable, ductile, metallic luster, solid at RT (except Hg), amphoteric
  • Nonmetal properties: Oxides are acidic, oxidizing agents, brittle, not malleable, not ductile, not metallic luster, state at RT can be solid, liquid, or gas
  • Atomic properties: Atomic number of elements: 118, Periods (Horizontal rows): 7, Groups/Family (Vertical columns): 18
• • Group A: Representative elements (s & p block), Group B: Transition elements (d block), Actinides & Lanthanides: Inner transition elements (f block)
• Periodic Trends:

  • Ionization energy: Energy needed to remove outermost electron in neutral atom
  • Electron affinity: Energy given off when neutral atom gains extra electron
  • Electronegativity: Ability of an atom to attract electron pair to itself, forming covalent bond
  • Atomic radius: ½ difference between nucleus of 2 atoms
  • Metallic property: Increase top to bottom, decrease left to right