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45 Cards in this Set
- Front
- Back
Arrhenius Acids
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Dissociate to produce an excess of hydrogen ions in solution (H+)
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Arrhenius Bases
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Dissociate to produce an excess of hydroxide ions in solution (OH-)
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Bronsted-Lowry Acids
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Species that can donate hydrogen ions
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Bronsted-Lowry Bases
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Species that can accept hydrogen ions
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Lewis Acids
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Species that accept electrons
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Lewis Bases
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Species that donate electrons in lone pairs
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Differences between Definitions
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Arrhenius is most restrictive and limited to aqueous solutions, Bronsted focuses on H, Lewis focuses on lone pairs
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Amphoteric Species
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Can behave as an acid or base
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Amphiprotic Species
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Amphoteric species that specifically behave as Bronsted-Lowry acids or bases
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Autoionization of Water
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Water Dissociation Constant
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Kw, is 10-14 at 298 K (25 C)
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pH
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pOH
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Relationship Between pH and pOH
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As pH increases, pOH decreases so that the sum is always 14 at 298 K
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Strong Acids and Basis
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Completely dissociate in solutions
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Common Strong Acids
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HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid), H2SO4 (sulfuric acid), HNO3 (nitric acid), and HClO4 (perchloric acid)
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Common Strong Bases
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NaOH (sodium hydroxide), KOH (potassium hydroxide) and other soluble hydroxides of group IA metals
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Converting Concentration to p Value
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[x] = 10-n then pX = n
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Estimate for Complicated Conversion to p Value
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Weak Acids and Bases
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Do not dissociate completely
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Impact of Autoionization with Strong Acids and Bases
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Only negligibale if concentration of acid or base is greater than 10^7
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Acid Dissociation Constant
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Base Dissociation Constant
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Dissociation Constant of Weak Acids and Basis
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The smaller the constant the weaker the acid or base, must be less than 1.0
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Strength of Conjugate
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Strong acids and bases have weak (inert) conjugates while weak acids and bases also have weak conjugates
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Induction
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Acids with electronegative elements near an acidic proton have increased acidic strength
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Neutralization Reactions
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React acids and bases together to form salts (and sometimes water)
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Combinations for Neutralization Reactions
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Resulting pH of Neutralized Solution
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Two strong: neutral, two weak: depends on relative strength of both, Strong acid: acidic, Strong base: basic
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Equivalents
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One mole of the species of interest (H+ for acids, OH- for bases)
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Normality
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The concentration of acid or base equivalents in solution equal to the multiplication of the concentration of the reactant by its number of equivalents
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Polyvalent Acids and Bases
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Can donate more than one equivalent
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Titrations
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Used to determine the concentration of a known reactant in a solution
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Titrant
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Has a known concentration and is added slowly to the titrand to reach the equivalence point
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Titrand
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Has an unknown concentration but a known volume
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Half Equivalence Point
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Midpoint of the buffering region in which half of the tritrant has been protonated or deprotonated
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Equivalence Point
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The steepest slope in a titration curve reached when the number of acid equivalents in the original solution equals the number of base equivalents added (or vice versa)
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Equivalence Point Trends
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Indicators
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Weak acids or bases that display different colors in their protonated and deprotonated forms
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Choosing an Indicator
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Should have a pKa close to the pH of the expected equivalence point
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Buffer Solutions
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Mixture of a weak acid and its conjugate salt or a weak base and its conjugate salt and are used to resist large fluctuations in pH
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Buffering Capacity
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Ability to resist changes in pH, maximum within 1 pH point of the pKa of the acid in the solution
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Henderson-Hasselbalch Equation- Weak Acid
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Henderson-Hasselbalch Equation- Weak Base
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IdentifyingType of Titration |
Identify starting point If pH >> 7: titrant is a strong base If pH > 7: titrant is a weak base If pH < 7: titrant is weak acid If pH << 7: titrant is a strong acid |