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25 Cards in this Set
- Front
- Back
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Bronsted Lowry Acid
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Substance that can donate protons (H+)
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B-L Base
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Substance that recieves and accepts protons
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Acid strength
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Linked to amount of H+ available
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H–A + H20 <--> H30+ A-
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Extent increases, Acidity Increases, K increases
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Factors that impact extent of inductive withdraw
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Examine electronegativity
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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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OH-, Oxides of 1A and 2A metals (except Mg, Be), H-, CH3-. Stronger the acid, weaker the conjugate base
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pH scale
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pH = -log[H3O+] or -log[OH-]
pH+pOH = 14 [H3O+] = 10-pH |
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Ionic compounds in competition
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Ammonium Acetate [NH4 C2H3O2]
Evaluate cation vs anion |
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Who wins competition?
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Evaluate extent of reaction
- Ka vs Kb |
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Examine reaction w/ H2O
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NH4 + H20 <--> NH3 + H30
Ka = [H3O][NH3] / [NH4] Kw/Kb |
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Examine Base in H20
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Kb'=Kw/Ka'
Compare Ka to Kb' |
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Ka> Kb'
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acidic
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Ka=kb'
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neutral
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Ka < kb'
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basic
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amphiprotic
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ions that can be acids or bases
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ion product of water
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The equilibrium constant for the self ionization of water.
2H20 <---> H3O+ + OH- Kw is the product of molar concentrations of hydronium ion and hydroxide ion. At STP, Kw= 1x10^-14 |
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Base dissociation constant (Kb)
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An equilibrium constant for the dissociation of a weak base in water
B+H20 <--> BH+ OH- kb=[BH+][OH-]/[B] |
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Buffer solution
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Solution that tends to resist pH change. Consists of a weak acid and its conjugate base. Or a weak base and its conjugate acid (NH3/NH4+)
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titration
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Analytical procedure where a solution of one reaction is slowly added to a soln of a second reactant. From the vol. and concentration of titrant, the second reactant can be calculated
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Standard solution
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any solution of accurately known concentration
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equivalence point
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The point at which amount of titrant added is exactly equal to amount of substance being titrated
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Endpoint
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Point at which indicator undergoes visible changes
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pH interval
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pH range over which the given acid-base indicator changes completely from one color to another.
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Equation involving Kb, Ka, Kw
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Ka = Kw / Kb
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