Unit 6 vocabulary : Acid/Base Equilibria Flashcard

Hydronium ion
The ion H3O, consisting of a protonated water molecule and present in all aqueous acids.
Bronsted-Lowry Acid
Any substance that donates hydrogen ion (proton).
Bronsted-Lowry Base
Any substance that accepts a proton.
Amphiprotic
Bronsted-Lowry acids and Lewis acids. It means a substance capable of donating and accepting protons, such as water and bicarbonate. Amphiprotic species must have acidic protons.
All amphiprotic substances are amphoteric; not all amphoteric species are amphiprotic.
Amphoteric
Lewis acids. It means a substance capable of acting both as a lone pair acceptor and a lone pair donor, such as aluminium hydroxide and tin dioxide. Amphoteric species don’t necessarily need acidic protons.
All amphiprotic substances are amphoteric; not all amphoteric species are amphiprotic.
Conjugate Base
Acids form a conjugate base.
Conjugate Acid
Bases form a conjugate acid.
Conjugate Acid-Base Pair
Ex) HNO3 + H2O -> H3O + NO3
A B CA CB
Autoionization
when H2O breaks down into H+ and OH-. the reaction is at equilibrium, because the H2O is always breaking down into H+ and OH-, which is constantly forming H2O.
It is at a pH 7 because the [H+] concentration is going back and forth at equilibrium, causing it to be neutral.
Ion Product Constant, Kw
KW = [H3O+][OH-]
pH
A measure of how acidic or basic (alkaline) an aqueous solution is.
Acid Dissociation Constant, Ka
The acid dissociation constant is the equilibrium constant of the dissociation reaction of an acid and is denoted by Ka.
Percent Ionization
The percentage of the weak electrolyte that ionizes in a solution of given concentration.
Polyprotic Acids
An Acid that can form two or more hydronium ions per molecule; often a least one step of ionization is weak.
Base Dissociation Constant, Kb
Equilibrium constant that measures the extent of dissociation for a base.
Hydrolysis
When water reacts with another substance and as a result the oxygen in water makes a bond with the substance.
Oxyacids
An oxyacid is an acid that contains an oxygen atom bonded to a hydrogen atom and at least one other element.
Ex) Sulfuric acid (H2SO4), phosphoric acid (H3PO4), and nitric acid (HNO3) are all oxyacids.
Carboxylic Acids
Carboxylic acid is an organic compound containing the COOH functional group.
Ex) acetic acid, CH3COOH
Lewis Acid
A substance that can accept an electron pair from a base; thus, AlCl3, BF3, and SO3 are acids.
Lewis Base
A substance that can donate an electron pair; examples are the hydroxide ion, OH-, and ammonia, NH3.
Common-Ion Effect
The lowering of the degree of ionization of a compound when another ionizable compound is added to a solution; the compound added has a common ion with the other compound.
Buffered Solutions (Buffers)
Buffer systems contain both an acid and a base(either an acid and its conjugate base or a base and its conjugate acid).
Buffer capacity
The relative ability of a buffer solution to resist pH change upon addition of an acid or a base.
Henderson-Hasselbalch equation
An equation expressing the pH of a buffer solution as a function of the concentration of the weak acid or base and the salt components of the buffer.
pH Titration Curve
A curve relating pH to the equivalents of strong base added per equivalent of acid in the solution.
Indicators
Indicators are dyes which change colors to show acidity or alkalinity. These are qualitative because they show acidity or alkalinity by changing color.

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