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66 Cards in this Set
- Front
- Back
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Vapor Pressure |
The pressure the gas exerts over the liquid at equilibrium. VP increases as T increases.
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Boiling Point
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The temperature at which the vapor pressure equals the ambient/environmental pressure
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Sublimation
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Solid → Gas
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Deposition |
Gas → Solid
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Triple Point
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Point at which the three phases exist in equilibrium
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Critical Point
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Where the phase boundary between the liquid and gas phases terminates. The temperature and pressure at which there is no distinction between the liquid and gas phases. The densities of the liquid and gas phases become equal. Heat of vaporization is 0 at T & P above the CP
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Multiple Component Phase Diagrams
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Phase diagram for a mixture of two or more components. Composition of the mixture, as well as T & P must be specified. Curves show the different compositions of the liquid phase and the vapor phase above a solution for different temperatures. Upper curve = composition of vapor. Lower curve = composition of liquid
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Colligative Properties
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Physical properties of solutions that are dependent on the concentration of dissolved particles, but not on their chemical identity. Vapor pressure depression, boiling point elevation, freezing point depression, osmotic pressure
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Water Phase Diagram
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Water has a solid/liquid equilibrium line with a negative slope
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Vapor Pressure Depression
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When you add solute to a solvent & it dissolves, the solvent in solution has a vapor pressure that is lower than the vapor pressure of the pure solvent for all temperatures
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Boiling Point Elevation
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The lowering of a solutions vapor pressure would mean a higher temperature is required to overcome atmospheric pressure thereby raising the boiling point (BP = the temp at which VP = AP)
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Van’t Hoff Factor
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The moles of particles dissolved into a solution per mole of solute molecules (i = 2 for NaCl
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Freezing Point Depression
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The presence of solute particles in solution interferes with the formation of the lattice arrangement of solvent molecules associated with the solid state. A greater amount of energy must be removed from solution (lower temperature) in order for it to solidify
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Osmotic Pressure
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Water pressure at which there is no net flow of water across a selectively permeable membrane with a solute dissolved in one side
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Solvation/Dissolution
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The electrostatic interaction between solute and solvent molecules (when water is the solvent → hydration). Involves breaking intermolecular attractions between solute molecules and between solvent molecules and forming new intermolecular interactions between solute and solvent
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Entropy
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A measure of molecular disorder, or the number of energy microstates available to a system at a given temperature (freedom to move around in different ways)
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General solubility rules for aqueous solutions
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1) All salts of alkali metals are water soluble 2) All salts of the ammonium ion (NH4+) are water soluble 3) All chlorides, bromides, and iodides are water soluble, with the exception of those formed with Ag+, Pb2+, and Hg22+ 4) All salts of the sulfate ion (SO42-) are water soluble, with the exceptions of those formed with Ca2+, Sr2+, Ba2+, and Pb2+ 5) All metal oxides are insoluble, with the exception of those formed with the alkali metals and CaO, SrO, BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides 6) All hydroxides are insoluble, with the exception of those formed with the alkali metals and Ca2+, Sr2+, and Ba2+ 7) All carbonates (CO32-), phosphates (PO43-), sulfides (S2-), and sulfites (SO32-) are insoluble, with the exception of those formed with the alkali metals and ammonium 8) All nitrate salts (NO3-) are water soluble
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Strong electrolyte
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a solute that dissociates completely into its constituent ions (NaCl, KCl, HCl in water)
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Weak electrolyte
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a solute that ionizes or hydrolyzes incompletely in aqueous solution and only some of the solute is dissolved into its ion constituents
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Nonelectrolytes
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compounds that do not ionize at all in aqueous solution (many nonpolar gases and organic compounds, such as O2, CO2, and glucose
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Percent Composition by Mass
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the mass of the solute divided by the mass of the solution (solute plus solvent), multiplied by 100%
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Mole Fraction
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The number of moles of the compound divided by the total number of moles of all species in the system. Used to calculate the vapor pressure depression of a solution and the partial pressures of gases in a system
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Molarity
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the number of moles of solute per liter of solution. Used for calculating the law of mass action, rate laws, osmotic pressure, pH and pOH, and the Nernst equation
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Molality
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the number of moles of solute per kilogram of solvent. For dilute aqueous solutions at 25 C, the molality is approximately equal to molarity, because the density of water at this temperature is 1 kg/L. Molality is required when calculating boiling point elevation and freezing point depression
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Normality
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The number of equivalents of solute per liter of solution. The concentration of the reactive species with which you are concerned (reaction-dependent)
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Gram equivalent weight/equivalent weight
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a measure of the reactive capacity of a molecule, the mass of a compound that produces one equivalent (1 mole of charge)
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Dilutions
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C1V1 = C2V2
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Solubility Product Constant (Ksp)
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The equilibrium constant for the solubility of a compound (don’t include pure solids or pure liquids in equilibrium expressions). Nongas solutes: Ksp increases with increasing T, Gas solutes: Ksp decreases with increasing T, increases with increasing P
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Ion Product
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Used to determine “where” the system is with respect to the equilibrium position. Concentrations = that of the ionic constituents at a given moment in time
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IP < Ksp
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solution is not yet at equilibrium (unsaturated)
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IP > Ksp
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solution is beyond equilibrium (supersaturated)
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Common Ion Effect
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the solubility of a salt in a solution is considerably reduced when it is dissolved in a solution that already contains one of its constituent ions (but has no effect on Ksp)/ Le Chatelier’s Principle
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Arrhenius acid
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A species that dissociates in water to produce a hydrogen ion, H+
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Arrhenius base
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A species that dissociates in water to produce a hydroxide ion, OH-
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Bronsted Lowry acid
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A species that donates hydrogen ions
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Bronsted Lowry base
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A species that accepts hydrogen ions
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Lewis acid
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An electron pair acceptor (Most inclusive definition: every Arrhenius acid is also a Bronsted-Lowry acid and every BL acid is also a Lewis acid)
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Lewis base
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An electron pair donor (Most inclusive definition: every Arrhenius base is also a BL base and every BL base is also a Lewis base)
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p-scale
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the negative logarithm of the number of items
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Strong acids
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HCl, HBr, HI, H2SO4, HNO3, HClO4
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Strong bases
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NaOH, KOH, other soluble hydroxides of Group IA & IIA
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Acid dissociation constant
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Ka, the smaller the Ka, the weaker the acid, the less it will dissociate
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Base dissociation constant
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Kb, the smaller the Kb, the weaker the base, the less it will dissociate
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Conjugate acid
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the acid formed when the base gains a proton
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Conjugate base
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the base formed when the acid loses a proton
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Amphoteric species
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acts like an acid in a basic environment & like a base in an acidic environment (H2O)
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Titration
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used to determine the molarity of a known reactant in a solution by reacting a known volume of a solution of unknown concentration with a known volume of a solution of known concentration
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Equivalence point
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when the number of acid equivalents present in the original solution equals the number of base equivalents added, or vice versa; approximate from graph: midpoint of region of curve with steepest slope
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Indicator
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weak organic acids or bases that have different colors in their protonated & deprotonated states, must be weaker than the acid/base being titrated, otherwise the indicator would be titrated first, want indicated to change color near the pH of the equivalence point
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End point
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point at which the indicator changes color; not equal to equivalence point, but volume difference is usually negligible
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Strong Acid/Strong Base Titration
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equivalence point pH = 7
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Weak Acid/Strong Base Titration
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equivalence point pH > 7
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Weak Base/Strong Acid Titration
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equivalence point pH < 7
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Polyvalent Acid/Base Titrations
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have multiple equivalence points
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Buffers
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mixture of a weak acid & its salt (which consists of its conjugate base and a cation) or a mixture of a weak base & its salt (which consists of its conjugate acid & an anion); resist changes in pH when small amounts of SA or SB are added; WA of buffer acts to neutralize the SB that is added (weak acid buffer solution) or WB of buffer acts to neutralize the SA that is added (weak base buffer solution)
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Rules for Assigning Oxidation Numbers
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1) The oxidation number of free elements is zero (N2, P4, S8, He) 2) The ON for a monoatomic ion = the charge of the ion 3) The ON of each group IA element in a compound is +1, each group IIA element is +2 4) The ON of each group VIIA element in a compound is -1 except when combined with an element of higher EN (HOCl – ON of Cl is +1) 5) The oxidation number of hydrogen is -1 in compounds with less electronegative elements than hydrogen (Groups IA and IIA). The more common oxidation number for hydrogen is +1 6) In most compounds, the oxidation number of oxygen is -2, except when combined with an element of higher electronegativity and in peroxides, [O-O]2- 7) The sum of the oxidation numbers of all the atoms present in a neutral compound is zero. The sum of the oxidation numbers of the atoms present in a polyatomic ion is equal to the charge of the ion
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Balancing Redox Reactions
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1) Separate the 2 half-reactions 2) Balance all atoms except H and O 3) In acidic solution, add H2O to balance the O atoms and then add H+ to balance the H atoms. In basic solution, add OH- and H2O to balance the O’s and H’s 4) Balance the charges of each half-reaction. Balance positive charge with electrons. The reduction half-reaction must consume the same number of electrons as are supplied by the oxidation half. 5) Add the half-reactions (cancel any H2Os, H+s, OH-s that appear on both sides of the equation) & confirm that mass and charge are balanced
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Salt Bridge
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Contains an inert electrolyte. Allows for the exchange of cations and anions
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Cell Diagram
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anode / anode solution // cathode solution / cathode
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Galvanic/Voltaic Cells
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Spontaneous redox reactions occur, -∆G, emf = + (sign on emf is always opposite the change in free energy), electrons flow from the anode (oxidation) to the cathode (reduction), anions flow from the salt bridge to the anode and cations flow from the salt bridge to the cathode
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Electrolytic Cells
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Nonspontaneous redox reactions occur, require the input of energy to proceed, +∆G, type of redox reaction driven by an external voltage source (Galvanic cell) = electrolysis, in which chemical compounds are decomposed. Can be used to drive nonspontaneous decomposition reactions. Oxidation occurs at the anode and reduction occurs at the cathode. Anions travel to the anode to be oxidized, cations travel to the cathode to be reduced.
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Faraday’s Constant (F)
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the amount of charge contained in one mole of electrons (1 F = 96, 487 C). On the MCAT you should round this number to 100,000 C/mol e-
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Concentration Cells
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Spontaneous redox reactions occur, -∆G, the electrodes are chemically identical, thus both have the same reduction potential. Therefore, current is generated as a function of a concentration gradient established between the two compartments resulting in a potential difference that drives the movement of electrons in the direction that results in equilibrium of the ion gradient. The current will stop when the concentrations of ion species in the half-cells are equal – emf = 0 when the concentrations are equal
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Reduction Potential
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Measured in volts (V), defined relative to the standard hydrogen electrode (0 V), the tendency of a species to gain electrons and be reduced, the more positive the reduction potential, the greater the species’ tendency to be reduced.
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Standard Reduction Potential
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Measured under standard conditions, 25 C, 1 M concentration for each ion participating in the reaction, a partial pressure of 1 atm for each gas participating in the reaction, and metals in their pure state
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Standard Electromotive Force
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the difference in potential between two half-cells under standard conditions. Determined by adding the standard reduction potential of the reduced species and the standard oxidation potential of the oxidized species. When adding standard potentials, do not multiply them by the number of moles oxidized or reduced |
