Acid and Bases (10) Flashcards
Arrhenius acid and base
Arrhenius acid will dissociate to form an excess of H+ in the solution.
The Arrhenius base will dissociate to form an excess of OH- in solution.
Bronsted-Lowry Acid
Bronsted-Lowry acid is a species that donates hydrogen ions (H+)
The bronsted-Lowry base is a species that accepts hydrogen ions (H+)
Water is a bronsted-lowry acid b/c it can donate a proton to other species.
Lewis acid-base
Lewis acid: electron pair acceptor
Lewis base: electron-pair donor
Bronsted lowry vs. Lewis acid base
Bronsted Lowry definition revolves around protons (H+)
Lewis definition of electrons
Acid naming:
If anion ends in -ite (less oxygen) , then acid will end with -ous acid
If anion ends in -ate (more oxygen), then acid will end with -ic acid
Water dissociation constant Kw
H2O + H2O ⇌ H3O+ + OH-
Kw = [H3O+] [OH-] = 1.0 x 10^-14
The equilibrium constant is dependent on temperature. Therefore, isolated changes in concentration, pressure, or volume will not affect Kw.
pH vs pOH
pH = -log [H+] = log 1/[H+]
pOH = -log[OH-] = log 1/[OH-]
pH + pOH = 14
pH condition
pH = 7 = neutral ONLY valid at 25°C (298K)
Strong acids and base
completely dissociate into their component ions in an aqueous solution.
NaOH –> Na+ + OH-
Strong acids
HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid), H2SO4 (sulfuric acid), HNO3 (nitric acid), and HClO4 (perchloric acid)
Strong base
NaOH (sodium hydroxide) , KOH (potassium hydroxide) , and other soluble hydroxides of Group IA metals.
Weak acids and base
partially dissociate in aqueous solution
Ka acid dissociation constant
the lower the Ka, the weaker the acid, the less it will dissociate
Kb base dissociation constant
the lower the Kb, the weaker the base, the less it will dissociate.
Kw, Ka, Kb
Ka x Kconjugate base = Kw
Kb x Kconjugate acid = Kw
Kw = 1.0 x 10^-14
acid equivalent vs. base equivalent
acid equivalent: equal to one mole of H+ (or more properly H3O+)
base equivalent: equal to one mole of OH-
Equivalence point
number of equivalents of acid and base are equal.
NaVa = NbVb
N = normality V = volume
pH of equivalence points
- Strong acid + weak base –> equivalence point pH < 7
- Strong acid + strong base –> equivalence point pH = 7
- Weak acid + strong base –> equivalence point pH > 7
Buffer System
weak acid + conjugate salt
weak base + conjugate salt
Buffer Capacity
The ability of a buffer to resist changes in pH. Maximum
buffering capacity is within 1 pH point of the pKa.
Half-equivalence point
center of the buffer region.
When [A-] = [HA] at the half equivalence point, log(1) = 0, so pH = pKa
Henderson- Hasselbalch equation
pH = pKa + log [A-] / [HA]
pOH = pKb + log [B+] / [BOH]
