AP Chem Final Exam Review

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Density
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mass/volume
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Law of Conservation of Mass
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mass reactants= mass products
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Law of Multiple Proportions
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atoms combine in fixed whole # ratios
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Alkali metals
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Group 1 metals
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Alkaline earth metals
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Group 2 metals
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Transition metals
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elements in groups 3-12
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mol
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6.022×10^23
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Limiting reactant
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reactant that’s completely used up in a chemical reaction
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excess reactant
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reactant which doesn’t get used up completely in a chemical reaction
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Theoretical yield
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amount of product produced when limiting reactant is used up
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experimental yield
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the actual amount of product produced in an experiment
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Molarity
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mol/L, concentration of a solution
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Molality
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mol/kg of solvent, used in calculating colligative properties
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soluble
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Group 1, Ammonium, Nitrates, Acetates, Sulfates, Halides
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insoluble
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Carbonates, Hydroxides, Oxides, Phosphates, Sulfides
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% yield
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experimental yield/theoretical yieldx100
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% error
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theoretical yield-experimental yield/theoretical yieldx100
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strong acids
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HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄
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strong bases
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Group 1 and heavier Group 2 bases
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strong acid strong base rxn
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H⁺+OH⁻→H₂O
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Strong acid weak base rxn
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H⁺+NH₃→NH₄
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weak acid strong base rxn
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HF+ OH⁻→H₂O
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standard solution
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a solution used in titrations whose concentration is known
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titrant
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buret solution used in titration
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analyte
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solution in flask being titrated
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equivalence point
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point where acid completely neutralizes base
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indicator
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a weak acid that changes color at or near the equivalence point
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end point
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point at which the titrated solution changes color
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oxidation
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loss of electrons, increase in oxidation #
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reduction
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gain of electrons, decrease in oxidation #
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oxidizing agent
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the reactant that is being reduced, brings about oxidation
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reduction agent
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the reactant that is being oxidized, brings about reduction
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Pressure
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force per unit area
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1 atm
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760 mmHg, 760 torr
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Standard Temperature and Pressure
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0.00°C, 1 atm
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22.4L
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volume of gas @STP
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Dalton’s Law
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Ptotal=Pa+Pb+Pc….
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mol Fraction
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mols A/ total mols, XA
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1/2mv²
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Kinetic Energy per molecule
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3/2RT
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Kinetic Energy per mol
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√3kT/m
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speed per molecule of gas
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√3RT/M(in kg)
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average speed of gas
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1.38×10⁻²³J/K
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Boltzmann constant, used in calculating speed of gas per molecule
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Graham’s Law
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effusion of a gas is inversely proportional to the square root of the molar mass
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r₁/r₂
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=√M₂/M₁
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3.0×10⁸m/s
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speed of light, C
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C
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=lambda(nu)
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lambda
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wavelength symbol
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nu
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frequency symbol
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E
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=h(nu)
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6.63×10⁻³⁴Js
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Planck’s constant, used to calculate energy w/frequency
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Heisenberg Uncertainty Principle
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we cannot simultaneously determine an atom’s exact path or location
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wavelength
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determined by the formula h/m(in kg)v, (v=velocity)
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Aufbau Principle
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e⁻s fill the lowest energy orbital first, then work their way up
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Pauli Exclusion Principle
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each orbital can hold two e⁻s each w/ opposite spins
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Hund’s Rule
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within a sublevel, place one e⁻ per orbital before pairing them
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linear
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AX₂, AX₂E₃
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trigonal planar
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AX₃
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tetrahedral
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AX₄
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trigonal pyramidal
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AX₃E
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bent
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AX₂E, AX₂E₂
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trigonal bipyramidal
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AX₅
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seesaw
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AX₄E
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t-shape
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AX₃E₂
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octahedral
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AX₆
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square pyramidal
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AX₅E
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square planar
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AX₄E₂
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single bond
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1 sigma bond
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double bond
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1 sigma bond, 1 pi bond
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triple bond
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1 sigma bond, 2 pi bonds
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system
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the part of the universe one is focused upon (in thermodynamics)
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surroundings
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the rest of the universe (in thermodynamics)
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endothermic
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positive enthalpy, heat flows into system
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exothermic
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negative enthalpy, heat flows into surroundings
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heat capacity
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heat required to raise the system 1°C
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Bond enthalpy
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ΔH when 1 mol of bonds is broken in the gaseous state
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work
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all forms of energy except for heat
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permanent gases
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substances w/ critical temperatures below 25°C
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boiling point
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point at which liquid→gas occurs
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melting point
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point at which solid→liquid occurs
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supercritical fluid
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substances above the critical temperature and pressure in which the pressure is so high that density and flowing ability of a \”gas\” resembles that of a liquid
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heat of vaporization
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energy required for liquid→gas
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heat of fusion
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energy required for melting to occur
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melting
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solid to liquid
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freezing
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liquid to solid
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vaporization
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liquid to gas
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condensation
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gas to liquid
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sublimation
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solid to gas
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deposition
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gas to solid
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Molecular substances
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-nonconductors of electricity – insoluble in water – low melting and boiling points
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London dispersion forces
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weakest IMFs, found in all molecules
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dipole dipole forces
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-electrostatic force between the positive and negative ends of a polar molecule -stronger than London forces, but weaker than H bonding
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Hydrogen bonding
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unusually strong dipole forces found when H is bonded to N, O, or F
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cohesion
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molecules’ tendency to stick to one another
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adhesion
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molecules’ tendency to stick to the container
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viscosity
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resistance to flow
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Network Covalent solids
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– atoms joined by a continuous network of covalent bonds – high melting points – nonconductors of electricity – insoluble in water – Ex: diamond, graphite, quartz
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Ionic solid
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– atoms joined by strong electrical forces between oppositely charged particles – high melting points – nonconducting as solids – conducting when molten or in solution – often water soluble
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metallic solids
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– metal cations held together by electrons not attached to any particular metal cation – high electrical conductivity – conduct heat – ductile and malleable -insoluble in water -high range of melting and boiling points
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Cg=kPg
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Henry’s Law, solubility of gases is directly proportional to the partial pressure of the gas
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P₁= X₁P₁°
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Raoult’s Law, relations between vapor pressure and concentrations
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Colligative properties
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vapor pressure lowering, boiling point elevation, freezing point depression, osmotic pressure
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ΔTb= kb x molality
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boiling point elevation formula
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ΔTf= kf x molality
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freezing point depression formula
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0.512°C
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kb of water
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1.86°C
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kf of water
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pi=(nRT)/v
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osmotic pressure formula
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Van’t Hoff factor
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for colligative properties for electrolytes, the # of mols of ions/ mols of solute
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rate
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how fast or slow a reaction occurs, becomes slower as it reaches equilibrium
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activation energy
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the minimum energy that molecules must possess for collisions to be effective, Ea
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catalyst
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lowers activation energy
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Q<K
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forward rxn occurs when
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Q>K
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reverse rxn occurs when
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LeChatelier’s Principle
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If a system @equilibrium is stressed, the system will shift so as to reestablish equilibrium
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Arrhenius acid
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puts H⁺ into solution
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Arrhenius base
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puts OH⁻ into solution
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Bronsted-Lowry acid
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proton donors
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Bronsted-Lowry base
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proton acceptors, must have an unshared pair of e⁻s
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conjugate acid
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the chemical formed when a base accepts a proton
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conjugate base
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the chemical formed when an acid donates a proton
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1.0×10⁻¹⁴
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Kw
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5% rule
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x can be ignored when % ionization is <5%
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Buffer
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a solution that resists a change in pH, contains both a weak acid and its conjugate base
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no precipitate forms
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if Q<Ksp
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a precipitate forms
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if Q>Ksp
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1st law of thermodynamics
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the total energy of the universe is constant, all systems tend towards minimum energy
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spontaneity
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the likelihood that a rxn will occur \”by itself\”
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entropy
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a measure of randomness or disorder
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2nd law of thermodynamics
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the total entropy is always increasing, all systems tend towards maximum entropy
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3rd law of thermodynamics
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the entropy of a pure perfectly formed crystal @0K is 0
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voltaic cells
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uses a spontaneous redox rxn to generate electrical energy, consists of 2 half cells
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cathode
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half cell in which reduction occurs
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anode
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half cell in which oxidation occurs
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salt bridge
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connects the 2 half cells in a voltaic cell
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96,500
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Faraday’s constant
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Nernst Equation
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Ecell= E°cell -RT/nF x lnQ
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flouride
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F¹⁻
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chloride
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Cl¹⁻
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iodide
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I¹⁻
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hydroxide
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OH¹⁻
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cyanide
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CN¹⁻
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chlorite
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ClO₂¹⁻
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chlorate
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ClO₃²⁻
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perchlorate
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ClO₄¹⁻
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bromate
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BrO₃¹⁻
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nitrite
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NO₂¹⁻
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nitrate
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NO₃¹⁻
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acetate
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C₂H₃O₂¹⁻
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permanganate
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MnO₄¹⁻
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carbonate
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CO₃²⁻
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chromate
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CrO₄²⁻
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dichromate
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Cr₂O₇²⁻
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oxalate
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C₂O₄²⁻
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oxide
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O²⁻
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sulfide
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S²⁻
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sulfite
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SO₃²⁻
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sulfate
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SO₄²⁻
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phosphate
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PO₄³⁻
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ammonium
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NH₄¹⁺

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