Concepts covered in Ch 16
Acid Base Definitions-Arrhenius/Bronsted Lowery.
Properties of Acids/BasesAutoionization of water and its equilibrium
Hydronium ion (another name for H+ in soln)
Monoprotic acid vs diprotic or triprotic acids
pH= - log [H+]
pOH= - log [OH-]
Strong acids/bases (100 % disassociates)
Calculating the pH of a SA/SB
Go straight from concentration to pH/pOH
Be careful for diprotic acids or bases with more
than l hydroxide.
Weak acids/bases (Less than 100% breaks apart and the rxn forms an equilibrium)
Strength of a WA determined by Ka values.
Bigger Ka values are stronger acids (with lower pH values)
Lower Ka values are weaker acids (with higher pH values)
Calculating pH of a WA
ID Conjugate acid/base pairs
Amphiprotic/amphoteric
Need its Ka + “ ICE”
Know the Properties of Acids/Bases.
Be able to write the equation for the autoionization of water.
Be able to calculate the H+/OH- and pH or pOH of any aqueous soln.
Kw = [H+][OH-]
@ 25 C Kw = 1 x 10 -14
pH = - log [H+]
pOH = - log [OH-]
pH + pOH = 14 @ 25 C
Remember that H+ and H3O+ are referring to the same thing when in an
aqueous environment.
Therefore you may see pH written as pH = - log [H3O+] or
[H3O+][OH-] = Kw
Know the Definitions of acids/bases.
Arrhenius:
Acids produce excess H+ in soln (or when dissolved in water)
Bases produce excess OH- in soln (or when dissolved in water)
Excess: more than what water possesses after it undergoes auto ionization.
According to this definition water can neither be an acid or a base. A condition for
acids/bases is that they be dissolved in water. This is the most restrictive definition.
Bronsted-Lowery:
Acids: donate a proton (H+)
Bases: accept a proton (H+)
This definition broadens the scope of acids and bases. First the substance does not
have to be an aqueous soln. Second, bases do not have to have hydroxide or produce
hydroxide in order for neutralization to occur.
According to this definition water can be both an acid or a base depending on what it is
doing
H2O + CO3-2 <--> HCO3 -1 +
OH acid
base
conj. acid
conj base
In the above situation carbonate is also a Arrhenius as well as a Bronsted Lowery
base, because it is a proton acceptor and in the process is creating excess OH- in soln.
The following situation shows an example of carbonate being only a Bronsted- Lowery
base.
HBr + CO3-2 <--> HCO3 -1 +Br acid
base
conj. acid
conj base
Notice that no excess OH- is being produced in the above rxn. Below is another
situation in which water is acting like a base. The ammonia is both Bronsted -Lowery
acid as well as a Arrhenius because it is producing excess H+ (hydronium ion)
Substances that in one situation are an acid and in another are a base are called
amphoteric.
H2O + NH4+1 <-->
base
acid
Acids
Conjugate base
H2CO3
HCO3-1
HClO
HCO3-1
CO3-2
ClO -1
Bases
Conjugate acid
NH3
S-2
HCO3-1
NH4 +1
HS-1
H2CO3
H3O +1 + NH3
conj acid
conj base
Be able to find the pH of a weak acid using ICE. Use short cut method first.
Step l.
Step 2.
Step 3
Step 4
Step 5.
Write the rxn for the disassociation of the acid and its Ka equation
Use ICE to calculate the amount of reactants/products @
equilibrium
State any assumptions (Assuming x is small……)
Calculate X by plugging into the Ka equation and solving.
Make sure that X is less than 5% of the original before calculating
pH. This is known as the percent that ionizes.
The final definition of acid or base is called the Lewis definition
Acids: electron pair acceptor (H+ is looking for a pair of electrons to make it
stable, but any other element or cpd will also be considered an acid if the need a
pair of electrons to create a stable octet)
Bases: electron pair donors. Any substance that has a nonbonded pair of
electrons in its outer shell can then act like a base and neutralize an H+ or other
cpd looking for a pair of electrons.
We did not have time to cover this last definition but it is the broadest definition and
encompasses more cpds that the first two definitions combined. Since these definitions
involved electrons, think Lewis dot. Or the electron definition.