BIOCHEMISTRY TOPICS

Acid-base chemistry: Introduction

Arrhenius definition. Acids: Strong and weak. Bases: Strong and weak. pH, pOH, and other logarithmic functions. The pH scale. The Brønsted-Lowry definition. Autoionization of water.

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Acids and bases are familiar to most of us, manly in terms of their general qualitative features, but a chemist seeks to understand them on a more fundamental level. An acquaintance with the chemistry of ions in solution serves as a good initial starting point for the study of acid-base chemistry, since acids and bases in aqueous solution (where is the solvent) can be viewed as the presence of specific types of ions. The first and simplest definition of an acid is the Arrhenius definition, which states that an acid is a substance that creates or consists of hydrogen ions in water, H⁺(aq). The corresponding Arrhenius definition of a base is a substance that creates or consists of hydroxide ions in water, OH^–(aq).

There are seven common strong acids (listed in the table below) that dissociate completely to H⁺(aq) and a conjugate anion in water. Other acids dissociate only partially - these are the weak acids. A primary example of a weak acid is the common organic acid acetic acid (ethanoic acid). Vinegar contains acetic acid at a low concentration. The formula for acetic acid is typically written as CH₃COOH, with the acidic hydrogen listed last. Examples of strong bases are the soluble metal hydroxides, the most common being sodium hydroxide (NaOH) and potassium hydroxide (KOH). Our primary example of a weak molecular base is ammonia, NH₃. When dissolved in water, a small proportion of the ammonia molecules react to form ammonium and hydroxide ions.

Acids and bases have long been known to be chemical opposites, as they undergo neutralization reactions. An acid-base titration is a quantitative, stoichiometric measurement of a neutralization reaction.

	Table 1 : Strong and weak acids
	Strong acids				Weak acids
	acid name and formula		type		acid name and formula		type
	hydrochloric acid - HCl hydrobromic acid - HBr hydroiodic acid - HI		hydrogen halides		HF- hydrofluoric acid		hydrogen halide
					phosphoric acid - H3PO4 sulfurous acid - H2SO3 chlorous acid - HClO2 hypochlorous acid - HClO carbonic acid - H2CO3		oxoacids
	nitric acid - HNO3 sulfuric acid - H2SO4 perchloric acid - HClO4 chloric acid - HClO3		oxoacids
							see special note below*
					formic acid - HCOOH acetic acid - CH3COOH benzoic acid - C6H5COOH		organic (carboxylic acids)
	*Special note on carbonic acid: H2CO3 decomposes to carbon dioxide and water: H2CO3(aq)   →  H2O(l) + CO2(aq) Carbon dioxide has limited solubility in water under atmospheric pressure, so the characteristic fizz of carbonated water is observed: CO2(aq)  →  CO2(g)

	Table 2 : Strong and weak bases
	Strong bases				Weak bases
	base name and formula		type		base name and formula		type
	lithium hydroxide - LiOH sodium hydroxide - NaOH potassium hydroxide - KOH rubidium hydroxide - RbOH cesium hydroxide - CsOH		alkali metal hydroxide		ammonia - NH3		molecular
					carbonate - CO32–(aq)		oxoanions
	tetraalkylammonium hydroxide, R4NOH #		quaternary ammonium hydroxide
					methylamine - CH3NH2		organic (amines, based on ammonia)
	sodium oxide - Na2O potassium oxide - K2O		metal oxides (base anhydrides)
					acetate - CH3COO–		organic anions (e.g. carboxylates)
	# Note: Like the metal hydroxides, these are compounds with ionic character. They are composed of the quaternary amine cation, tetraalkylammonium (R is an alkyl group such as methyl, CH3), and the hydroxide anion : [R4N+] and OH–, respectively. Some sources consider insoluble hydroxides such as Mg(OH)2 to be weak bases. They create relatively little OH–(aq) because of their low solubility.

pH, pOH, and other logarithmic functions

In order to determine not only whether a given compond acts as an acid or base when mixed with water, but also to measure how acidic or how basic a given aqueous solution may be, the pH scale is defined. The pH is a logarithmic function of the H⁺(aq) concentration,

pH  =  − log [H⁺].

Note the minus sign in the definition creates an inverse relationship between [H⁺] and pH. A very high [H⁺] means pH is low, and vice-versa. But pH varies much less than [H⁺] in magnitude. This is characteristic of logarithmic functions, and makes them useful in cases where quantities vary over many orders of magnitude. Another example of such a quantity is K, the equilibrium constant for a reaction. We will want to be able to use pK values, where

pK  =  − log K

to quantify the strength of acids on the same scale as the pH scale.

In the range of hydronium ion concentrations we will typically encounter, the pH ranges between 0 and 14. As we may well know already, a pH of 7 - that is, [H⁺] = 1.0 × 10⁻⁷ M - is considered neutral, and pH less than 7 is considered acidic, while pH > 7 is basic. As a measure of basicity, and a counterpart to pH, we introduce pOH, defined as

pOH  =  − log [ OH ⁻ ].

There is a simple but important relationship between pH and pOH in aqueous solutions of acidic and basic compounds

pH  +  pOH  =  14

which has its basis in the definition and value of the ionization constant K_w, which is discussed further below.

The Brønsted-Lowry definition

The Arrhenius definitions for acid and base are completely dependent on the premise of an aqueous solvent system. Although we are typically working with aqueous solutions, the Arrhenius definition turns out to be just too limited. If an acid-base reaction occurs in aqueous solution, surely water can and does act as a participant. Even if the reaction entails no net consumption or production of water, water molecules can act as intermediaries. But is this necessarily the case? In many cases, direct transfer of H⁺ from an acid to a base (other than hydroxide) occurs, and the Arrhenius definition does not cover this type of reaction. To deal with this and to free ourselves from having to be limited to aqueous systems, the Brønsted-Lowry approach is used. The Brønsted-Lowry definition of an acid and a base, which simply states that an acid is a proton (H⁺) donor, and a base is a proton acceptor. The Brønsted-Lowry definition is quite general, and we find it convenient to consider a general Brønsted-Lowry acid, HA, and its corresponding acid dissociation equation:

HA  =  H⁺  +  A^–

Here, HA is the acid, and A^– is the conjugate base of HA. Brønsted-Lowry acids and bases always occur in conjugate pairs, and one of the merits of the acid dissociation equation is that it helps make this fact clear. As a specific example, water can act as an acid when it loses a proton to form hydroxide ion, the conjugate base of water:

H₂O  =  H⁺  +  OH ^– .

It is important for us to realize at the outset that water and some other substances can act as either acids or bases. Such substances are termed amphiprotic. If water is a base, then its conjugate acid must be the result of water accepting a proton. This species, H₃O⁺, is called hydronium ion. The acid dissociation equation for hydronium ion is:

H₃O⁺  =  H⁺  +  H₂O.

One point that is often confusing for students learning acid-base chemistry is that "H⁺(aq)" is often used as a shorthand for hydronium ion, or any species representing a proton transferred to water. What we really mean in that case is H⁺(aq), and we should take special care to keep a distinction in mind between H⁺(aq) and H⁺ as written in the acid dissociation equation for any Brønsted acid. In an acid dissociation equation, H⁺ is meant to represent a proton being transferred to another species acting as a base, and thus the equation only represents half of a complete Brønsted-Lowry acid-base reaction. For example, the formation of hydrochloric acid in water is the result of the combination of two acid dissociation equations:

HCl  =  H⁺  +  Cl ^– .

H₃O⁺  =  H⁺  +  H₂O.

The way to combine these is to reverse one of the half-reactions and add the two equations together so that there is an acid and a base on each side of the equation, and the H⁺'s cancel out. The summed equation is thus a proton transfer reaction, which is another way of saying "Brønsted-Lowry acid-base reaction". But which acid dissociation equation should we flip? Well, it really doesn't matter - either way we will get a valid chemical equation from the sum. But if we want to write the summed equation so that the products are favored at equilibrium, we should flip the acid dissociation equation for the weaker acid. We will learn how to quantify acid strength in short order, but for now let's utilize the fact that HCl is a stronger acid than hydronium ion, so let's flip the bottom equation and then sum the equations:

HCl  =  H⁺  +  Cl ^– .

H⁺  +  H₂O  =  H₃O⁺.

sum:  HCl  +  H₂O   =   H₃O⁺  +  Cl ^–.

We have written a proton-transfer (or acid-base) reaction in a form that favors the products at equilibrium. It is worth noting and remembering that the products favored at equilibrium in an acid-base reaction are invariably the weaker acid and the weaker base.

Autoionization of water

What if we put together the two acid dissociation equations for water as an acid and hydronium ion?

H₃O⁺  =  H⁺  +  H₂O.

H₂O  =  H⁺  +  OH ^– .

Following the rules above, we would write

H₃O⁺  =   H⁺  +  H₂O

H⁺  +  OH ^–   =   H₂O

sum:  H₃O⁺  +  OH ^–  =  H₂O  +  H₂O.

You may recognize the summed equation as the net ionic equation for the neutralization reaction between a strong acid and a strong base. The reverse of this reaction - written with the products extremely unfavorable at equilibrium - is the so-called autoionization of water. The equivalent terms autodissociation and autoprotolysis are used in some texts.

H₂O  +  H₂O   =   H₃O⁺  +  OH ^–.

We can quantify water's small tendency to autodissociate by defining a special type of equilibrium constant called K_w and assigning an experimentally determined value to it.

K_w = [ H₃O⁺ ][ OH ^– ] = 1.0 × 10^–14   (at 25°C).

Note that this is the equilibrium constant expression that we would write for the chemical equation representing the autoionization equation for water (considering the concentration of water as a constant). The very small value for K_w reflects the very weak propensity of water to autoionize.

Definition of K_a, the acid dissociation constant

For the general acid dissociation equation,

HA  =  H⁺  +  A^–

we define K_a, the acid dissociation constant, by forming the expression

K_a  =  [H⁺][A^– ] / [HA].

Note how we apply the normal rules for writing an equilibrium constant - concentrations of products in the numerator, concentrations of reactants in the denominator - to the acid dissociation equation in defining K_a. We will use K_a as a quantitative measure of acid strength.

More about acid-base chemistry...

Related topics pages:

• weak acid / weak base equilibria
• polyprotic systems
• buffers
• titrations
• water