Chapter 16 (Acids/bases) Flashcards
Acid-Base theory
Define Bronsted-Lowry acids and bases.
Acid: H+ donor
Base: H+ acceptor
Define Arrhenius acids and bases
Acid: H+ producer
Base: OH- producer
Define Lewis acids and bases
Acids: e- acceptor
Bases: e- donor
Name the 7 strong acids
HCl, HNO3, H2SO4, HClO, HI, HBr, and HClO3
Name the 7 strong bases
KOH, NaOH, Ba(OH)2, CsOH, Sr(OH)2, LiOH, and Ca(OH)2
What is the conjugate acid of NH3?
NH4+ ; NH3 accepts a proton.
What is the conjugate base of CH3COOH?
CH3COO- ; CH3COOH lost a proton.
Define Kw
1*10^-14
This is the ionization constant of water.
Given that [NaOH] = 0.1 M, is it possible to calculate the pH without an ICE table?
Yes; simply for strong acids and bases that ionize completely, the concentration of the base is equal to the concentration of OH- or H+ (respectively). We are assuming total disassociation.
When does one use an ICE table?
Use an ice table to find either unknown concentrations of H+/OH-, or to find the value of Ka given the concentrations.
Define a polyprotic acid.
This type of acid has more than one H+ to donate in solution. This means they have a Ka value for each proton that comes off.
How many Ka constants does H3PO4 have?
3
How do bond strength and acid strength relate?
These two are inversely related. The stronger the bond, the weaker the acid.
Define a weak acid.
Weak acids do not completely deionize (fall apart) in water, which is why Ka is needed to tell us how much of the substance will dissociate in solution.
How do anion size and acid strength relate?
The larger the anion, the stronger the acid. There is less overlap in the orbitals between the proton and the anion, that it is able to fall off easier in solution.
