Definitions of Acids and Bases


Bronsted-Lowry Definition of Acids and Bases

We will use the Bronsted-Lowry definitions for acids and bases:

Acids are species that donate a proton (H+).
and
bases are species that accept a proton.

Acid example:
HNO3 (aq) + H2O <--> NO3-(aq) + H3O+(aq)

Keq = a very large number

In this example, HNO3 is an acid and H2O is acting as a base.
NO3- is called the conjugate base of the acid HNO3, and H3O+ is the conjugate acid of the base H2O.

Base example:
NH3 (aq) + H2O <--> NH4+(aq) + OH-(aq)

K = 1.8x10-5

In this example, NH3 is a base and H2O is acting as an acid. NH4+ is the conjugate acid of the base NH3, and OH- is the conjugate base of the acid H2O.

A compound that can act as either an acid or a base, such as the H2O in the above examples, is called amphiprotic.


Kw, pH, and pOH

H2O + H2O <--> H3O+(aq) + OH-(aq)
or
H2O <--> H+(aq) + OH-(aq)

Kw = [H3O+][OH-] = [H+][OH-] = 1.00x10-14 (at 25oC)

(Using [H3O+] is equivalent to using [H+].)

Kw is called the dissociation constant or ionization constant of water.

In pure water [H+] = [OH-] = 1.00x10-7 M.

pH is a shorthand notation for -log[H+]
and
pOH is a shorthand notation for -log[OH-].

pH + pOH = 14.

Solutions are called
neutral when pH = 7, [H+] = [OH-] = 1.00x10-7
acidic when pH < 7, [H+] > 1.00x10-7
basic when pH > 7, [H+] < 1.00x10-7

Example: What is the pH of a solution of 0.025 M HNO3? (See example 13.4 in text.)

HNO3 is a strong acid and for all practical purposes dissociates completely.

HNO3(aq) + H2O <--> NO3-(aq) + H3O+(aq)

[H+] = 0.025 M

pH = -log(0.025 M) = 1.6

What is the pOH of this solution? There are 2 ways to calculate pOH:

Kw = 1.00x10-14 = [0.025 M][OH-]
[OH-] = 4.00x10-13
pOH = -log(4.00x10-13) = 12.40

or:

pOH = 14.00 - pH = 14.0 - 1.60 = 12.40

What are pH and pOH for a 0.0025 M solution of HNO3?

pH = -log(0.0025 M) = 2.60
pOH = 14.00 - 2.60 = 11.40

Notice that pH and pOH change by 1 for a factor of 10 change in [H+] and [OH-].


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