Ch. 10: Acids and Bases Flashcards
Arrhenius Acids
Dissociate to produce an excess of hydrogen ions in soln
Arrhenius Bases
Dissociate to produce and 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 acceptors
Lewis Bases
Electron-pair donors
Statements
All Arrhenius acids and bases are Bronsted-Lowry acids and bases, and all bronsted-lowry acids and bases are lewis acids ands and bases; however, the converse of these statements is not necesarily true
Amphoteric
Species which can behave as acids or bases
Amphiprotic
Species which can specifically behave as Bronsted lowry acids or bases
Water
Classic example of an amphoteric, amphiprotic species– it can accept a hydrogen ion to become a hydronium ion, or it can donate a hydrogen ion to become a hydroxide ion
Conjugate Species of Polyvalent acids and bases
Can also behave as amphoteric and amphiprotic species
Water Dissociation constant Kw
10^-14 at 298K. Like other equilibrium constants, Kw is only affected by changes in temp
pH and pOH
Can be calculated given the concentrations of H3O+ and OH- ions, respectively. In aqueous solns, pH+pOH=14 at 298K
Strong acids and bases
Completely dissociate in solution
Weak acids and bases
Do not completely dissociate in solution and have corresponding dissociation complexes (Ka and Kb)
Conjugate Acids and Bases
In the Bronsted-Lowry def, acids have conjugate bases that are formed when the acid is deprotonated. Bases have conjugate acids that are formed when the base is protonated
- Strong acids and bases have very weak (inert) conjugates
- Weak acids and bases have weak conjugates
Neutralization reactions
From salts and (sometimes) water
Equivalent
Defined as one mole of the species of interest
Normality
In acid-base chemistry, normality is the concentration of acid or base equivalents in solution
Polyvalent
Acids and bases that can donate or accept multiple electrons. The 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.
Titrations
Used to determine the concentration of a known reactant in a soln
Titrant
Has a known concentration and is added slowly to the titrand to reach the equivalence point
Titrand
Has an unknown concentration but a known volume
Half-equivalence point
Midpoint of the buffering 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 eqivalents in the original soln equals the number of base equivalents added or vice versa
pH values of acids and bases
- Strong acid and strong base titrations have equivalence points at pH = 7
- Weak acid and strong base titrations have equivalencce points at pH > 7
- Weak base and strong acid titrations have equivalencce points at pH < 7
- Weak acid and weak base titrations can have equivalence points above or below 7 depending on the relative strength of the acid and base
Indicator
Chosen for a titration, should have a pKa close to the pH of the expected equivalence point
Endpoint
When the indicator reaches its final color in a titration
Polyvalent acid/base titrations
Multiple buffering regions and equivalence points are observed
Buffer solutions
Consist of a mixture of a weak acid and its conjugate salt or a weak base 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 soln
Henderson-Hasselbalch Equation
Quantifies the relationship between pH and pKa for weak acids and between pOh and pKb for weak bases; when a soln is optimally buffered, pH = pKa and pOH = pKb
