Acids and Bases Flashcards
bronsted acid
- donates H+
lewis acid
- accepts e-
- electrophile
bronsted base
- accepts H+
lewis base
- donates e-
- nucleophile
acid
- e- acceptor
- H+ donor
- electrophile
- leaving group
- ox. agent (gets reduced)
- cation
base
- e- donor
- H+ acceptor
- nucleophile
- part that isn’t LG
- reducing agents (gets oxidized)
- anion
- ligand/chelate
- coordinate bond
recognizing acids
- generally have more electronegative atoms bonded to H
recognizing bases
- generally have less electronegative atoms with lone pairs
atoms without H
- can be acids if electron deficient or with large positive charges
atoms without lone pairs
- are not usually basic
amphoteric compounds
- have characteristics of both acids and bases
stability of conjugate
- tells strength of original
- if very stable, is not very strong so the original is strong
increases acidity
- more positive charge
- more electronegative atom
- larger atom
increases basicity
- more negative charge
- less electronegative
- smaller atom
acid dissociation constant
[A:-][H3O+]
_________ = Ka
[HA}
increased Ka
- increased numerator
- increased products
- increased acidity
decreased Ka
- decreased acidity
base dissociation constant
[BH+][HO-]
_________ = Kb
[B}
increased Kb
- increased numerator
- increased products
- increased basicity
decreased Kb
- decreased basicity
strong acids
- dissociate completely
- Ka > 1
- equilibrium favors products
- produce bases so weak they’re not basic
weak acids
- partially dissociate
- Ka < 1
- equilibrium factors reactants
- produce strong conjugate bases (not necessarily strong bases)
strong bases
- dissociate completely
- Kb > 1
weak bases
- partially dissociate
- Kb < 1
common strong acids
- H2SO4 (diprotic)
- HClO4
- HCl
- HBr
- HI
common strong bases
- O2- (diprotic)
- OH-
- OR-
- NH2-
- NR2-
- H-
- R-
acidic salt
- contains an ion that is a weak acid
basic salt
- contains an ion that is a weak base
group 1 and 2 cation
- are not acidic
what does p mean
- inverse
formula for pH
- pH= -log[H+]
formula for pKa
- pKa= -log[Ka]
autoionization of H2O formulas
pH + pOH = 14
KaKb = Kw = 1e-14
pKa + pKb = PKw = 14
as A increases
- Ka increases
- pKa decreases
as B increases
- Kb increases
- pKb decreases
strong acid or base calculation
- assume complete dissociation
- pH = - log [H3O+]
weak acid or base calculation
- use Ka or Kb to determine how much dissociates
Ka = [A:-](H3O+]
______________
[HA]
buffer calculation
- use Henderson Hasselbach
buffers
- mixtures of conjugate acid/base pairs
- minimize changes in pH
Henderson-Hasselbach equation
pH = pKa + log [A-]/[HA]
a reaction between an acid and a base
- neutralization reaction
conjugate acid/base of a strong acid/base
- unreactive
- spectator ion
- pH neutral
conjugate acid/base of a weak acid/base
- weakly reactive but still somewhat reactive
half equivalence point
- where it goes more horizontal
equivalence point
- mol H+ = mol OH-
- goes vertical again
where is the equivalence point in a titration of a strong base with a weak acid?
- greater than 7
where is the equivalence point in a titration of a weak base with a strong acid
- less than 7
titration curve of a diprotic acid with a strong base
- 2 half equivalence and 2 equivalence points
- first equivalence point will be less than 7. The second equivalence point will be greater than 7
color indicators in terms of pH and pKa
- pH < pKa color 1
- pH = pKA mix of color 1 and color 2
- pH > pKa color 2
way to remember Lewis acid/base
- same lewis from dot structures concerned with electrons
algebraic solution from ICE table
pH = -1/2 log (Ka [WA])
OR
pOH = -1/2 log (Kb [WB])
neutralization reactants are always
- exothermic
half equivalence point
- point at which 50% of the acid is dissociated.
- pKa of the weak acid or base
- if you have the base then 14-pKa=pKb
first equivalence point of a diprotic
- 1/2 (pKa1 + pKa2)
- also called isoelectric point
