Final study guide of concepts
Final study guide of concepts Chemistry 1220
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This 9 page Study Guide was uploaded by fehlman.1 Notetaker on Wednesday April 20, 2016. The Study Guide belongs to Chemistry 1220 at Ohio State University taught by Dr. Loza in Winter 2016. Since its upload, it has received 100 views. For similar materials see Chemistry 1220 in Chemistry at Ohio State University.
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Date Created: 04/20/16
Chemistry Final Exam Concept Study Guide: Chapter 13: Solutions can either be -Miscible: components mix and dissolve completely -Immiscible: components form multiple layers - Solutions must have similar intermolecular forces to mix Delta H solution = Delta H1 + Delta H2 + Delta H3 - Delta H1 = energy required to separate solution (always positive) (the smaller this value, the more readily the solution occurs) - Delta H2 = energy required to separate solvent (always positive) (In many cases, this value is negligible because it is so small) - Delta H3 = energy gained from solute-solvent mixing (Hydration energy) Solution formation: -Physical process: the solute has the ability to be recovered unchanged -Chemical process: The solute is changed, cannot be recovered to its original state Like dissolves like: Nonpolar solutes with dissolve in a nonpolar solution, polar solutes will dissolve in a poplar solution Solutions can be: -Saturated: additional solute added will not dissolved -Unsaturated: less solute than the equilibrium amount of solute -Supersaturated: Greater amount of solute than needed to form a saturated solution Pressure changes only affect gases, the relationship between these gases and solubility is Henry’s Law: S = KP, solubility of gas = constant(1.38x10^- 3)X pressure of gas Temperatures effects on solubility - In an endothermic reaction: addition of heat will increase solubility - In an exothermic reaction: addition of heat will decrease solubility Colligative properties: What happens when additional solute is added to a solution? - Lowering of vapor pressure - Boiling point elevation - Freezing point lowering - Osmotic pressure Electrolytes have free floating ions in water, Non-electrolytes do not Vapor Pressure lowering - Raoult’s law: XaPa = Psoln; (mole fraction of solvent)(Vapor pressure of pure solvent) = (vapor pressure of the solution) - Vapor pressure of multiple liquids = Dalton’s Law: Ptotal = Pa + Pb - When vapors separate, the one with the higher pressure will leave first Boiling point elevation: - Increases with addition of solute - Delta Tb = Kbm: Boiling point elevation = molal BP elevation constant X molality of the solute Freezing point lowering - Freezing point is lowered as solute is added - Delta Tf = Kfm; Freezing point lowering = molal FP depression constant X molality of solute Osmotic Pressure - Pressure required to prevent osmotic diffusion - Osmotic pressure = M(molarity) X R(0.0821 atm/mol) X T(temperature in Kelvin) Colloids: Intermediate between solutes and solutions - Scatters light via the Tyndall effect - Can be Hydrophilic: water loving or Hydrophobic: water hating - Colloids can be removed by either adding heat or a salt Chapter 14: Factors that effect reaction rates: - Concentration of reactants - Surface area of reactants - Temperature of reaction - Presence of a catalyst Rate of reaction = change in moles/ change in time -Stoichiometric coefficients play a role in rate relationships. Rate Law Equation: Rate = k[A]^x[B]^y - K is the rate constant - X and y are the order of each reaction. Reaction order = sum of the exponents in the rate law Zero Order Reactions: - Integrated rate law equation: [A]t = -akt + [A]0 - Half life equation: T1/2 = [A]0/2ak First Order Reactions: - Integrated rate law: ln[A]t-ln[A]0 = -akt - Half life equation: T1/2 = 0.693/ak Second order: - 1/[A]t = akt + 1/[A]0 - Half life equation: T1/2 = 1/ak[A]0 The rate constant of a reaction will increase as temperature is increased Activation energy(Ea): minimum amount of energy required for a chemical reaction - Ea = Potential energy of activated complex – potential energy of reactants - Delta E of the reaction = Potential energy of the products – potential energy of the reactants Arrhenius Equation: depends on temperature - K = Ae^-Ea/RT - K= rate constant - A is a constant - Ea = Activation energy (kj) - R = 8.314 J/mol k - T = Temperature (K) The rate determining step in a reaction mechanism is the slowest step Unimolecular- single molecule Bimolecular- 2 molecules collide Termolecular- 3 molecules collide Catalyst: Provides alternate reaction pathway with a lower Ea - Speeds up forward and reverse reactions by the same amount 4 step process of a reaction: - Adsorption - Activation - Reaction - Desorption Ezymes: large lock and key catalysts in the human body Chapter 15: Equilibrium: concentrations of all products and reactants do not change with time. Forward rate = reverse rate K (equilibrium constant) = Products^x/Reactants^y Solid and liquid states are left out of the equilibrium constant expression When K is greater than 1 equilibrium lies to the right (products favored) When K equals about 1, the reverse and forward reactions are approximately equal When the direction of the chemical reaction is reversed, the reciprocal of k is taken When the chemical reaction is multiplied or divided by a coefficient, k is raised to that exponent Converting from Kc to Kp: - Kc = Kp(RT)^-delta n Kp = Kc(RT)^delta n Reaction Quotient Q, is used when concentration values are substituted into the [Products] and [Reactants] If… - Q = K, reaction is at equilibrium - Q is greater than K, reverse reaction predominates - Q is less than K, forward reaction predominates Le Chatliers Principle: a system in equilibrium is disrupted by the addition or subtraction of something. - If something is added to the left side of the equation, equilibrium shifts to the right and vice versa - If something is taken away from the left side of the equation, equilibrium shifts to the left As pressure increases, volume decreases and vice versa Changing temperature will affect K Chapter 16: Acids and Bases Strong acids: HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO4 Strong bases: group I & II hydroxides Strong acids and bases are single arrowed dissociations Weak acids and bases have a double arrow for dissociation and K is less then 1 Arrhenius acids and bases: - Uses H2O as the solvent - Revolves around the production of OH (base) and H (acid) Bronsted Lowry acids and bases: - Acid: donates a proton - Base: gains a proton - Amphoteric- species can act as an acid or a base - Conjugate acids and bases Lewis acids and bases: - Acid: accepts an electron pair - Base: donates an electron pair Water auto ionizes itself and produces a kw of 1.0 X 10^-14, this is used to switch between Ka and Kb of reactions - Kw/Kb = Ka, Kw/Ka = Kb pH and pOH pH = -log[H+] POH = -log[OH-] [H+] = 10^-pH [OH-] = 10^-pOH pH + POH = 14 Ka and Kb expressions are the same as K: products over reactants Pka = -logKa Polyprotic acids: have more than 1 ionizable proton per molecule A salt breaks apart and forms with H+ and OH- ions, from this, the new formations are analyzed to determine whether the new formations are basic or acidic, determining then if the salt is basic or acidic Chapter 17: Occurs when a weak acid salt is added to a weak acid, and there is a common ion - ICE table is restructured as more ion is added - As % ionization decreases, pH increases when a common weak acid salt is added to a weak acid solution. Buffer resist change in pH because the contain acid and basic species to neutralize H+ and OH- - Buffer capacity: the amount of acid/base a buffer can neutralize before the pH begins to change - The higher the molarity of the buffer, the greater the buffer to resist pH change Henderson Hasselbach Equation: pH = pKa + log ([A-]/[HA]) Acid Base titrations - Equivalence point: moles of acid and base are equivalent - End point: indicator changes color 4 titration zones - Initial concentrations - Excess acid - Equivalence point - Excess base Solubility and solubility equilibrium - Ksp = Products/Reactants - Smaller Ksp = insoluble salt - Larger Ksp = soluble salt Factors that effect solubilty: - Temperature - Presence of other solutes Solubility of a slightly soluble salt is decreased by the presence of a second solute that funishes the common ion Solubility of a compound containing a basic anion increases as the solution becomes more acidic Formation of complex ions: Metal ion + lewis base - Salt solubility increases if the metal ion could form a complex ion. Amphoteric Hydroxides and oxides - Can act as acid or base - Amphoteric: Al+3, Cr+3, Zn+2, Sn+2 - Not amphoteric: Ca+2, Fe+2, Fe+3, only react with acids K = Kf x Ksp Solubility decreases as concentration increases pH goes up as molar solubility increases By calculating Q and comparing it to Ksp - If Q is larger than Ksp, precipitation, shift left - If Q = Ksp, reaction is at equilibrium - If Q is less than Ksp, no precipitation occurs, shift right The smaller ion concentration always precipitates first in a 1:1 ratio chemical reaction. Chapter 19: First Law of thermodynamics: energy is conserved and neither created or destroyed State functions = change in heat, enthalpy, and entropy Gibbs Free Energy: - Gibbs free energy = enthalpy – Temperature x entropy - If delta G is negative, the reaction is spontaneous - If delta G = 0, the reaction is at equilibrium - If delta G is negative, the reaction is not spontaneous Spontaneous vs nonspontaneous reactions Second Law of thermodynamics: an increase in the disorder of a gas is a increase in the entropy - Reversible process: Change in entropy of the universe is 0 - Irreversible process: change in entropy of the universe is greater than 0 When a chemical reaction goes from more atoms to less atoms, there in a decrease in entropy Atoms have degrees of freedom: ways they can have motion and energy - Translate: move in 1 direction - Vibrate: periodically toward and away - Rotate: spin like a top Third law of thermodynamics: if temperature increases, so does motion of particles ie. Disorder When concentrations show the system is not at equilibrium a different Gibbs equation is used - Delta G = Delta G not + RTlnQ - R = 8.314 J/mol K - T = absolute temp. (K) - Q = reaction quotient Chapter 20: Electrochemistry: Oxidation: loss of electrons, increase in oxidation # - Oxidizing agent: gain electron, oxidize other substances, oxidation agents are always reduced Reduction: Gain of electrons, decrease in oxidation # - Reducing agent: loss electrons, reduce other substances, reducing agent always oxidized Disproportionation: an element is both oxidized and reduced Number of atoms and charge must be the same on both sides of the equation. Voltaic Cell: - Reduction (+): cathode - Oxidation (-): anode - A porous barrier needs to be present to complete the electrical circuit Half reactions are combined with reduction potentials in order to calculate the cell potential Negative delta G says the reaction is spontaneous Equation: delta G = -nFE - N = number of overall electrons transferred - F = Faradays constant (96,500J) Cell EMF and equilibrium: E = E not – 0.0592/n log Q In an electrolytic cell: - The anode is positive - The cathode is negative Chapter 23: Coordination sphere = Metal ion + ligand Donor atom: binds to central metal atom Unidentate: bond through only one donor atom Polydentate: bind simultaneously through multiple donor atoms Cations are named as they are in salts, with the cations being named before the anions Ligands are named in alphabetical order, then the metal - Unidentate: DI,tri, tetra, penta, hexa, - Polydentate: Bis, tris, tetrakis Anionic Ligands end in suffix O - Ide = o - Ate = ato - Ite = ito Neutral ligand names are the same as the molecule with the exception of H2O (aqua) and NH3 (ammine) Naming the metal: - Cation or neutral = same as metal - Anion = suffix ate - Oxidation # is the roman numeral following the name of the metal Isomerism: 2 or more compounds of the same composition but different arrangement of atoms. - Linkage: unidentate has 2 different atoms for coordination - Coordination: ligand exchange between cation and anion - Mirror image ligands don’t super impose - Stereoisomers: same bonds different special arrangements - Geometric isomers Electrons go from lower, to higher d orbitals Crystal field theory: - Bond between metal and ligand is ionic - All orbitals in free metal ions have the same energy - Ligands are arranged in order of magnitude of orbital splitting - Cl-, F-, H2O, NH3, en, NO2-,CN-
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