General Chemistry Ch 10. Acids and Bases Flashcards
Arrhenius acids
Dissociate to produce an excess of hydrogen ions in solution
Arrhenius bases
Dissociate to produce an excess of hydroxide ions in solution
Bronsted-Lowry acids
Species that can donate hydrogen ions
Bronsted-Lowry bases
Species that can accept hydrogen ions
Lewis acids
Electron pair accepts
Lewis bases
Electron pair donors
Arrhenius/Bronsted Lowry connection
All Arrhenius acids/bases are Bronsted-Lowry acid/bases but reverse not necessarily true
Bronsted-Lowry/Lewis connection
All Bronsted-Lowry acids/bases are Lewis acids/bases but reverse not necessarily true
Amphoteric species
Those that can behave as an acid or base, water good example, also conjugate species of polyvalent acids/bases
Amphiportic species
Amphoteric species that specifically can behave as a Bronsted-Lowry acid or Bronsted-Lowry base, water good example, also conjugate species of polyvalent acids/bases
Water dissociation constant
Kw = 10^-14 at 298 K, only affected by changes in temperature
pH and pOH
Can be calculated given the concentrations of H3O+ and OH- ions, respectively, in aqueous solutions pH+pOH=14
Strong acids/bases
Completely dissociate in solution, very weak/inert conjugates
Weak acids/bases
Do not completely dissociate in solution and have corresponding dissociation constants (Ka and Kb respectively), weak conjugates
Conjugate bases
Formed when a Bronsted-Lowry acid is deprotonated
Conjugate acids
Formed when a Bronsted-Lowry base is protonated
Neutralization reactions
Form salt and water
Equivalent
One mole of the species of interest
Normality
In acid base chemistry, the concentration of acid or base equivalents in solution
Polyvalent
Acids and bases that can donate or accept multiple electrons, normality of a solution containing a polyvalent species is the molarity of the acid or base times the number of protons it can donate or accept, multiple offering regions and equivalence points observed during titration
Titrations
Used to determine the concentration of a known reactant in solution
Titrant
Has a known concentration and is added slowly to the tetrad to reach the equivalence point during a titration
Titrand
Has an unknown concentration but a known volume during a titration
Half equivalence point
The midpoint of the buggering region, in which half of the titrant has been protonated or deprotonated, thus [HA] = [A-] and a buffer is formed
Equivalence point
Indicated by the steepest slope in a titration curve, it is reached when the number of acid equivalents in the original solution equals the number of base equivalents added or vice versa
Strong acid and strong base titration
Equivalence point is around pH=7
Weak acid and strong base titration
Equivalence point at pH>7
Weak base and strong acid titration
Equivalence point at pH<7
Weak acid and weak base titration
Can have equivalence points above or below 7 depending on the relative strength of the acid and base
Indicators
Weak acids or bases that displace different colors in their protonated and deprotonated forms, the one chosen for titration should have a pKa close to the pH of the expected equivalence point, endpoint of a titration is when the indicator reaches its final color
Buffer solutions
Consist of a mixture of a weak acid and its conjugate salt or a weak base and its conjugate salt, they resist large fluctuations in pH
Buffering capacity
Refers to the ability of a buffer to resist changes in pH, maximal buffering capacity is seen within 1 pH point of the pKa of the acid in the buffer solution
Henderson-Hasselbalch equation
Quantifies the relationship between pH and pKa for weak acids and between pOH and kPb for weak bases, when a solution is optimally buffered, pH=pKa and pOH=pKb
