Acids and Bases

What are acids and bases? There are three major definitions. We will look at two in which the proton is a major defining component (the third, Lewis definition, is not needed for our study).

(Although I will signify protons in water as H⁺, you should realize that naked protons do not exist in water - they are always hydrated. At a minimum we see the hydronium ion, H₃O⁺. But hydronium ion is in fact also generally thought to be hydrated, so you will sometimes see hydrogen ion represented as H₅O₂⁺, H₇O₃⁺, etc.)

• Arrhenius Definition:

• Acids release protons (H⁺) into water.
• Bases release hydroxide ions (OH^-) into water.
• This is very limited - it only deals with acids and bases in water, and many substances which chemists and others think of as bases (such as ammonia) don't fit the definition. Thus we will focus on the Brønsted-Lowry definition (Brønsted definition in abbreviation):

• Brønsted Definition:

• Acids are proton donors.
• Bases are proton acceptors.
• Note that there is no restriction as to solvent, and many substances besides hydroxide ion can contribute basicity.
• A consequence of the Brønsted definition is that all acids and bases are related to one or more conjugate bases or acids. That is, when an acid dissociates to give a proton, it also generates a conjugate base which can react with (accept) a proton in the reverse reaction. For example, in the case of water:

				H2O		H+	+	OH-
				acid				conj. base
		H+	+	OH-		H2O
				base		conj. acid
H3O+		H+	+	H2O		OH-	+	H+
conj. acid				acid base		conj. base

Strong vs. Weak Acids & Bases

These terms have nothing to do with concentration, rather they refer to the degree of dissociation of an acid or base:

• A Strong acid is 100% dissociated at all concentrations up to 1M. Common strong acids include:
  • Nitric acid (HNO₃)
  • Hydrochloric acid (HCl)
  • Sulfuric acid (H₂SO₄) for the first dissociation only: H₂SO₄ HSO₄^- + H⁺. The second dissociation is weak, that is it hardly dissociates at 1M.
• A Weak acid is only partly dissociated at 1M. The degree of dissociation varies widely, from a few percent to an infinitesimal degree. Common weak acids include:
  • Acetic acid (HC₂H₃O₂ or CH₃CO₂H, etc.)
  • Formic acid (HCO₂H)
  • Hydrofluoric acid (HF)
  • Most acids of biological origin such as amino acids, fatty acids, metabolites, nucleic acids etc.
• A Strong base is 100% dissociated at all concentrations up to 1M. Common strong bases include:
  • Sodium hydroxide (NaOH)
  • Potassium hydroxide (KOH)
• A Weak base is only partly dissociated at 1M. The degree of dissociation varies widely, from a few percent to an infinitesimal degree. Common weak acids include:
  • Ammonia (NH₃)
  • Aluminum hydroxide (Al(OH)₃)

Aqueous Hydrogen Ion Concentration and pH

The concentration of hydronium ion in water is extremely influential on all kinds of chemistry. The range of hydronium ion concentration in water is also vast, with extremes of about 10M to about 10^-15M, and commonly ranging from 1M - 10^-14M. Imagine plotting [H₃O⁺] vs. volume of acid added to a base solution in a titration. If you had one cm on the graph paper = 10^-14M, then you would need a piece of paper 10⁹ km long (greater than the distance from the Sun to Jupiter) to plot this titration! Obviously a more convenient measure is needed. This is easily accomplished by looking instead at the logarithm of [H⁺] and defining a new term,

Turns out that the concentration of hydrogen ion in water is related to the concentration of hydroxide ion due to the equilibrium dissociation of water:

H₂O H⁺ + OH^-, so

pH

Let's look at pH a bit:

• Low pH means acidic:
  • For 1M strong acid, pH = 0 (log 1 = 0)
  • For 0.1M strong acid, pH = 1
  • For 10^-7M H⁺, pH = 7
• High pH means basic:
  • For 1M strong base, pH = 14 ([H⁺] = (1.0 x 10^-14) / [OH^-] = (1.0 x 10^-14) / 1 = 1.0 x 10^-14 and pH = -log(1.0 x 10^-14) = 14.
  • For 0.0010M [OH^-], pH = 11.00 (Note significant figures: the digits to the left of the decimal are the power of 10, so are NOT significant, only the digits to the right are significant. Thus the 2 s.f. concentration, 1.0 x 10^-3 gives .00 for 2 s.f.

Examples:

• What is the pH of a solution of 0.015 M HCl?  

• What is the pH of a solution of 0.067 M NaOH

Note that the "p" has the more general meaning of "-log[]". Thus pOH is -log [OH^-], pCa = -log [Ca²⁺], etc.

C107 Laboratory

C107 Home

C107 Lecture Notes

© R A Paselk

Last modified 3 December 2009