Chapters 16, 17, and 19 – Flashcards
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| strong acids |
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| HCl, H2SO4, HI, HBr, H3O, HClO3, HClO4, HNO3 |
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| strong bases |
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| NaOH, KOH, Ca(OH)2, O2, ionic hydroxides of alkali metals, heavy alkaline earth metals |
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| Bronsted-Lowry Acid |
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| proton donator |
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| Bronsted-Lowry Base |
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| proton acceptor |
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| amphiprotic |
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| can function as a Bronsted-Lowry acid or base depending on what it reacts with |
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| conjugate base |
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| species that remains when a proton is removed from the acid |
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| conjugate acid |
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| species that is formed by adding a proton to the base |
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| autoionization |
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| water ionizes slightly to from H+ and OH- |
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| ion-product constant |
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| Kw = [H+] [OH-] = 1.0e-14 at 25 degrees C |
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| pH |
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| -log[H+] or {14 - (-log[OH-])} |
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| acid dissociation constant |
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| Ka, representative of weak acids, equilibrium constant for (HA) --> (H+) + (A-) |
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| percent ionization |
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| (concentration ionized / original concentration) x 100 ( [H+] equilibrium) / ( [HA] initial) x 100 |
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| polyprotic acids |
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| more than one ionizable proton (Ka1, Ka2...) |
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| base dissociation constant |
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| Kb, representative of weak bases (amines, NH3...), equilibrium constant for the reaction B + H2O --> HB+ + OH- |
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| Kw = Ka x Kb |
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| hydrolysis |
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| the reaction of ions with water resulting in a pH change |
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| do not undergo hydrolysis |
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| the cations of the alkali metals and alkaline earth metals, anions of strong acids (always spectator ions in acid-base chemistry) |
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| tendency of a substance to show acid or base characteristics is dependent upon: |
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| bond polarity, bond strength, X- ion stability |
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| Lewis Acid |
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| electron pair acceptor |
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| Lewis base |
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| electron pair donator |
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| pH + pOH = 14 |
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| Ka |
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| [H+][A-] / [HA] |
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| Kb |
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| [HB+][OH-] / [B] |
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| common-ion effect |
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| the dissociation of a weak acid or weak base is repressed by the presence of a strong electrolyte that provides an ion common to the equilibrium |
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| pH = pKa + log (base / acid) |
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| Ksp |
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| equilibrium constant that expresses the extent to which the compound disolves, equal to the product of the concentration of the ions involved in the equilibrium each raised to the power of its coefficient |
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| molar solubility |
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| number of moles of the solute that dissolve in forming a liter of a saturated solution of the solute (mol/L) |
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| solubility |
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| grams of solute dissolved in L of solution |
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| reversible process |
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| results in no change in entropy (S=0) |
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| irreversible process |
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| results in an increase in overall entropy (S>0) |
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| the number of microstates available increases with an increase in: |
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| volume, temperature, number of molescules |
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| increase in entropy |
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| increase in temperature, volume or number of particles |
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| if G is negative: |
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| the reaction is spontaneous |
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| if G is 0: |
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| equilibrium |
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| if G is positive |
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| nonspontaneous and work must be applied |
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| isothermal process |
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| any process that occurs at a constant temperature |
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| at a constant temperature, the entropy of a system is: |
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| given by the heat absorbed by the system along a reversible path (Qrev) divided by the temperature |
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| W |
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| number of microstates present |
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| standard molar entropy for an isothermal process is |
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| equal to -∆H / T |
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| when ∆G is negative |
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| the process is spontaneous |
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| when ∆G is positive |
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| the process is nonspontaneous |
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| at equilibrium ∆G is |
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| equal to 0 and the process is reversible |
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| entropy term |
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| -T∆S |
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| melting of ice |
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| ∆H>0, ∆S>0, nonspontaneous at low temperatures, and spontaneous at high temperatures |
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| at equilibrium Q=K |
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| the standard free energy change is directly related to the equilibrium constant for the reaction |
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| calculating [H+] given [OH-] |
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| [H+] = (1.0e-14 / [OH-]) |
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| calculating the pH of a strong acid |
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| [H+] = [ion], so pH = -log [ion] |
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| calculating the pH of a strong base |
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| [OH-] = [strong base] so pOH = 14 - (-log [strong base]) |
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| calculating Ka from measured pH |
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| 10^(measured pH) = [H+], ICE tablewith [H+] as change, Ka = ([products] / [reactants]) |
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| using Ka to calculate pH |
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| ionization equation, ICE table where change is "x", Ka equation using "x", solve for -log(x) |
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| using Ka to calculate percent ionization |
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| ICE table using x as change, Ka equation solve for x, if greater than 5% of molar value use the quadratic equation to find [ion], then use [ion] / [original] x100 |
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| using pH to determine salt concentration |
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| pOH = 14 - pH, [OH-] = 10^pOH, ICE table with x as [initial] and change is [OH-], Kb equation solving for x |
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| calculating Ka given Kb |
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| Ka = (1e-14) / Kb |
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| if common ion is involved in calculating [ion] |
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| x is small relative to other concentrations |
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| calculating pH of a buffer |
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| pH = pKa + log ([base] / [acid]) |
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| preparing a buffer |
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| pOH = 14 - pH, [OH-] = 10^pOH, Kb equation solving for [other ion], then [other ion] x volume of solution |
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| calculating pH for strong acid-strong base titration |
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| find moles of H+ and OH-, ICE table subtracting mol OH- from mol H+, final mol H+ / sum of volumes = [H+] |
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| calculate pH for weak acid-strong base titration |
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| find moles of acid and base, subtract mol base from mol acid and add mol base to conjugate base, final mol acid / sum of volumes, final mol conjugate base / sum of volumes, Ka equation solving for other concentration |
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| calculating pH at equivalence point |
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| find mol acid then divide by (volume x2), ICE table with x for change, Kb equation solving for x to find [OH-], then change that to pH |
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| calculating Ksp from Solubility |
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| write balance ionic equation, use coefficient multiplication to solve for moles not given, Ksp = [products] / [reactants] no solids, or pure liquids |
