63 Acids and Bases

Nature of Acids and Bases

Acids and bases will neutralize one another to form liquid water and a salt.

Acids

Acids have long been recognized as a distinctive class of compounds whose aqueous solutions exhibit the following properties:

• A characteristic sour taste.
• Changes the color of litmus from blue to red.
• Reacts with certain metals to produce gaseous [latex]\text{H}_2[/latex].
• Reacts with bases to form a salt and water.

Acidic solutions have a [latex]\text{pH}[/latex] less than 7, with lower [latex]\text{pH}[/latex] values corresponding to increasing acidity. Common examples of acids include acetic acid (in vinegar), sulfuric acid (used in car batteries), and tartaric acid (used in baking).

There are three common definitions for acids:

• Arrhenius acid: any substances that increases the concentration of hydronium ions ([latex]\text{H}_3\text{O}^+[/latex]) in solution.
• Brønsted-Lowry acid: any substance that can act as a proton donor.
• Lewis acid: any substance that can accept a pair of electrons.

Acid Strength and Strong Acids

The strength of an acid refers to how readily an acid will lose or donate a proton, oftentimes in solution. A stronger acid more readily ionizes, or dissociates, in a solution than a weaker acid. The six common strong acids are:

• hydrochloric acid ([latex]\text{HCI}[/latex])
• hydrobromic acid ([latex]\text{HBr}[/latex])
• hydroiodic acid ([latex]\text{HI}[/latex])
• sulfuric acid ([latex]\text{H}_2\text{SO}_4[/latex]; only the first proton is considered strongly acidic)
• nitric acid ([latex]\text{HNO}_3[/latex])
• perchloric acid ([latex]\text{HClO}_4[/latex])

Each of these acids ionize essentially 100% in solution. By definition, a strong acid is one that completely dissociates in water; in other words, one mole of the generic strong acid, [latex]\text{HA}[/latex], will yield one mole of [latex]\text{H}^+[/latex], one mole of the conjugate base, [latex]\text{A}^-[/latex], with none of the unprotonated acid [latex]\text{HA}[/latex] remaining in solution. By contrast, however, a weak acid, being less willing to donate its proton, will only partially dissociate in solution. At equilibrium, both the acid and the conjugate base will be present, along with a significant amount of the undissociated species, [latex]\text{HA}[/latex] .

Factors Affecting Acid Strength

Two key factors contribute to overall strength of an acid:

• polarity of the molecule
• strength of the H-A bond

These two factors are actually related. The more polar the molecule, the more the electron density within the molecule will be drawn away from the proton. The greater the partial positive charge on the proton, the weaker the H-A bond will be, and the more readily the proton will dissociate in solution.

Acid strengths are also often discussed in terms of the stability of the conjugate base. Stronger acids have a larger [latex]\text{K}_a[/latex] and a more negative [latex]\text{pK}_a[/latex] than weaker acids.

Bases

There are three common definitions of bases:

• Arrhenius base: any compound that donates an hydroxide ion ([latex]\text{OH}^-[/latex]) in solution.
• Brønsted-Lowry base: any compound capable of accepting a proton.
• Lewis base: any compound capable of donating an electron pair.

In water, basic solutions will have a [latex]\text{pH}[/latex] between 7-14.

Base Strength and Strong Bases

A strong base is the converse of a strong acid; whereas an acid is considered strong if it can readily donate protons, a base is considered strong if it can readily deprotonate (i.e, remove an [latex]\text{H}^+[/latex] ion) from other compounds. As with acids, we often talk of basic aqueous solutions in water, and the species being deprotonated is often water itself. The general reaction looks like:[latex]\text{A}^- (aq) + \text{H}_2\text{O} (aq) \rightarrow \text{AH} (aq) + \text{OH}^- (aq)[/latex]

Thus, deprotonated water yields hydroxide ions, which is no surprise. The concentration of hydroxide ions increases as pH increases.

Most alkali metal and some alkaline earth metal hydroxides are strong bases in solution. These include:

• sodium hydroxide ([latex]\text{NaOH}[/latex])
• potassium hydroxide ([latex]\text{KOH}[/latex])
• lithium hydroxide ([latex]\text{LiOH}[/latex])
• rubidium hydroxide ([latex]\text{RbOH}[/latex])
• cesium hydroxide ([latex]\text{CsOH}[/latex])
• calcium hydroxide ([latex]\text{Ca(OH)}_2[/latex])
• barium hydroxide ([latex]\text{Ba(OH)}_2[/latex])
• strontium hydroxide ([latex]\text{Sr(OH)}_2[/latex])

The alkali metal hydroxides dissociate completely in solution. The alkaline earth metal hydroxides are less soluble but are still considered to be strong bases.

Acid/Base Neutralization

Acids and bases react with one another to yield water and a salt. For instance:

[latex]\text{HCl} (aq) + \text{NaOH} (aq) \rightarrow \text{H}_2\text{O} (l) + \text{NaCl} (aq)[/latex]

This reaction is called a neutralization reaction.

Acids + Bases Made Easy! Part 1 – What the Heck is an Acid or Base? – Organic Chemistry – YouTube: Ever wondered what the heck an Acid or Base actually is? Were you ever super confused in high school or college chemistry? In this video I introduce to you guys what the heck an Acid and Base really is forgetting the Lewis or Bronstead/Lowry definitions and then we’ll go more in depth in parts 2,3, and 4.

The Arrhenius Definition

An Arrhenius acid dissociates in water to form hydrogen ions, while an Arrhenius base dissociates in water to form hydroxide ions.

The Arrhenius Definition

An acid-base reaction is a chemical reaction that occurs between an acid and a base. Several concepts exist that provide alternative definitions for the reaction mechanisms involved and their application in solving related problems. Despite several differences in definitions, their importance as different methods of analysis becomes apparent when they are applied to acid-base reactions for gaseous or liquid species, or when acid or base character may be somewhat less apparent.

The Arrhenius definition of acid-base reactions, which was devised by Svante Arrhenius, is a development of the hydrogen theory of acids. It was used to provide a modern definition of acids and bases, and followed from Arrhenius’s work with Friedrich Wilhelm Ostwald in establishing the presence of ions in aqueous solution in 1884. This led to Arrhenius receiving the Nobel Prize in Chemistry in 1903.

As defined by Arrhenius:

• An Arrhenius acid is a substance that dissociates in water to form hydrogen ions ([latex]\text{H}^+[/latex]). In other words, an acid increases the concentration of [latex]\text{H}^+[/latex] ions in an aqueous solution. This protonation of water yields the hydronium ion ([latex]\text{H}_3\text{O}^+[/latex]); in modern times, [latex]\text{H}^+[/latex] is used as a shorthand for [latex]\text{H}_3\text{O}^+[/latex] because it is now known that a bare proton ([latex]\text{H}^+[/latex]) does not exist as a free species in aqueous solution.
• An Arrhenius base is a substance that dissociates in water to form hydroxide ([latex]\text{OH}^-[/latex]) ions. In other words, a base increases the concentration of OH^– ions in an aqueous solution.

Limitations of the Arrhenius Definition

The Arrhenius definitions of acidity and alkalinity are restricted to aqueous solutions and refer to the concentration of the solvated ions. Under this definition, pure [latex]\text{H}_2\text{SO}_4[/latex] or [latex]\text{HCl}[/latex] dissolved in toluene are not acidic, despite the fact that both of these acids will donate a proton to toluene. In addition, under the Arrhenius definition, a solution of sodium amide ([latex]\text{NaNH}_2[/latex]) in liquid ammonia is not alkaline, despite the fact that the amide ion ([latex]\text{NH}_2^-[/latex]) will readily deprotonate ammonia. Thus, the Arrhenius definition can only describe acids and bases in an aqueous environment.

Arrhenius Acid-Base Reaction

An Arrhenius acid-base reaction is defined as the reaction of a proton and an hydroxide ion to form water:

[latex]\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}[/latex]

Thus, an Arrhenius acid base reaction is simply a neutralization reaction.

Chemistry 12.1 What are Acids and Bases? (Part 1 of 2) – YouTube: This introduction to acids and bases discusses their general properties and explains the Arrhenius definitions for acids and bases.

The Brønsted-Lowry Definition of Acids and Bases

A Brønsted-Lowry acid is any species capable of donating a proton; a Brønsted-Lowry base is any species capable of accepting a proton.

Originally, acids and bases were defined by Svante Arrhenius. His original definition stated that acids were compounds that increased the concentration of hydrogen ions ([latex]\text{H}^+[/latex]) in solution, whereas bases were compounds that increased the concentration of hydroxide ions ([latex]\text{OH}^-[/latex]) in solutions. Problems arise with this conceptualization because Arrhenius’s definition is limited to aqueous solutions, referring to the solvation of aqueous ions, and is therefore not inclusive of acids dissolved in organic solvents. To solve this problem, Johannes Nicolaus Brønsted and Thomas Martin Lowry, in 1923, both independently proposed an alternative definition of acids and bases. In this newer system, Brønsted-Lowry acids were defined as any molecule or ion that is capable of donating a hydrogen cation (proton, [latex]\text{H}^+[/latex] ), whereas a Brønsted-Lowry base is a species with the ability to gain, or accept, a hydrogen cation. A wide range of compounds can be classified in the Brønsted-Lowry framework: mineral acids and derivatives such as sulfonates, carboxylic acids, amines, carbon acids, and many more.

Brønsted-Lowry Acid/Base Reaction

Keep in mind that acids and bases must always react in pairs. This is because if a compound is to behave as an acid, donating its proton, then there must necessarily be a base present to accept that proton. The general scheme for a Brønsted-Lowry acid/base reaction can be visualized in the form:

[latex]\text{acid + base } \rightleftharpoons \text{ conjugate base + conjugate acid}[/latex]

Here, a conjugate base is the species that is left over after the Brønsted acid donates its proton. The conjugate acid is the species that is formed when the Brønsted base accepts a proton from the Brønsted acid. Therefore, according to the Brønsted-Lowry definition, an acid-base reaction is one in which a conjugate base and a conjugate acid are formed (note how this is different from the Arrhenius definition of an acid-base reaction, which is limited to the reaction of [latex]\text{H}^+[/latex] with [latex]\text{OH}^-[/latex] to produce water). Lastly, note that the reaction can proceed in either the forward or the backward direction; in each case, the acid donates a proton to the base.

Consider the reaction between acetic acid and water:

[latex]\text{H}_3\text{CCOOH} (aq) + \text{H}_2\text{O} (l) \rightleftharpoons \text{H}_3\text{CCOO}^- + \text{H}_3\text{O}^+ (aq)[/latex]

Here, acetic acid acts as a Brønsted-Lowry acid, donating a proton to water, which acts as the Brønsted-Lowry base. The products include the acetate ion, which is the conjugate base formed in the reaction, as well as hydronium ion, which is the conjugate acid formed.

Note that water is amphoteric; depending on the circumstances, it can act as either an acid or a base, either donating or accepting a proton. For instance, in the presence of ammonia, water will donate a proton and act as a Brønsted-Lowry acid:

[latex]\text{NH}_3 (aq) + \text{H}_2\text{O} (l) \rightleftharpoons \text{NH}_4^+ (aq) + \text{OH}^- (aq)[/latex]

Here, ammonia is the Brønsted-Lowry base. The conjugate acid formed in the reaction is the ammonium ion, and the conjugate base formed is hydroxide.

Chemistry 12.1 What are Acids and Bases? (Part 2 of 2) – YouTube: This lesson continues to describe acids and bases according to their definition. We first look at the Brønsted-Lowry theory, and then describe Lewis acids and bases according to the Lewis Theory

Acid-Base Properties of Water

Water is capable of acting as either an acid or a base and can undergo self-ionization.

Under standard conditions, water will self-ionize to a very small extent. The self-ionization of water refers to the reaction in which a water molecule donates one of its protons to a neighboring water molecule, either in pure water or in aqueous solution. The result is the formation of a hydroxide ion ([latex]\text{OH}^-[/latex]) and a hydronium ion ([latex]\text{H}_3\text{O}^+[/latex]). The reaction can be written as follows:

[latex]\text{H}_2\text{O} + \text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{OH}^-[/latex]

This is an example of autoprotolysis (meaning “self-protonating”) and it exemplifies the amphoteric nature of water (ability to act as both an acid and a base ).

pH, pOH, and Other p Scales

A p-scale is a negative logarithmic scale.

pH and pOH

Recall the reaction for the autoionization of water:

[latex]\text{H}_2\text{O} \rightleftharpoons \text{H}^+ (aq) + \text{OH}^- (aq)[/latex]

If we take the negative logarithm of each concentration at equilibrium, we get:

[latex]\text{pH} = \text{-log} [ \text{H}^+ ] = \text{-log} \left( 1.0 \times 10^{-7} \right) = 7.0[/latex]

[latex]\text{pOH} = \text{-log} [ \text{OH}^- ] = \text{-log} \left( 1.0 \times 10^{-7} \right) = 7.0[/latex]

Here we have the reason that neutral water has a pH of 7.0 -; this is the pH at which the concentrations of [latex]\text{H}^+[/latex] and [latex]\text{OH}^-[/latex] are exactly equal.

Lastly, we should take note of the following relationship:

[latex]\text{pH + pOH = 14}[/latex]

This relationship will always apply to aqueous solutions. It is a quick and convenient way to find pH from [latex]\text{pOH}[/latex], hydrogen ion concentration from hydroxide ion concentration, and more.

In acid-base chemistry, the amount by which an acid or base dissociates to form [latex]\text{H}^+[/latex] or [latex]\text{OH}_-[/latex] ions in solution can be given by the raw concentration values. However, because these values are often very small for weak acids and weak bases, the p-scale is used to simplify these numbers and make them more convenient to work with.

Interactive: pH: Test the pH of things like coffee, spit, and soap to determine whether each is acidic, basic, or neutral. Visualize the relative number of hydroxide ions and hydronium ions in solution. Switch between logarithmic and linear scales. Investigate whether changing the volume or diluting with water affects the pH. Or you can design your own liquid!

pH and pOH: This lesson introduces the pH scale and discusses the relationship between pH, [latex]\text{H+}[/latex], ([latex]\text{OH-}[/latex]) and [latex]\text{pOH}[/latex].