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21.1: Acid-Base Definitions

Created by: CK-12

Lesson Objectives

• Describe the properties of acids and bases.
• Define an acid and a base according to the Arrhenius theory.
• Define an acid and a base according to the Brønsted-Lowry theory. Be able to identify the conjugate acid-base pairs in a Brønsted-Lowry acid-base reaction.
• Define an acid and a base according to the Lewis theory.

Lesson Vocabulary

• amphoteric
• Arrhenius acid
• Arrhenius base
• Brønsted-Lowry acid
• Brønsted-Lowry base
• conjugate acid
• conjugate acid-base pair
• conjugate base
• hydronium ion
• Lewis acid
• Lewis base
• monoprotic acid
• polyprotic acid

Recalling Prior Knowledge

• What are the rules for naming acids?
• How are bases named?

Acids and bases are classes of compounds that you most likely encounter every day. Common household items such as fruits, juices, soaps, and detergents all contain various acids and bases. In this lesson, you will learn about the properties of acids and bases and several ways to define exactly what is considered an acid or base.

Acid-Base Properties

Acids are very common in some of the foods that we eat. Citrus fruits such as oranges and lemons contain citric acid and ascorbic acid, which is better known as vitamin C (Figure below (A)). Carbonated sodas contain phosphoric acid. Vinegar contains acetic acid. Your own stomach utilizes hydrochloric acid to digest food. Bases are less common as foods, but they are nonetheless present in many household products (Figure below (B)). Many cleaners contain ammonia, a base. Sodium hydroxide is found in drain cleaner. Antacids, which combat excess stomach acid, are comprised of bases such as magnesium hydroxide or sodium hydrogen carbonate.

(A) Lemons and other citrus fruits contain citric and ascorbic acids. (B) Ammonia is a base that is present in many household cleaners.

Acids

Acids are a distinct class of compounds because of the properties of their aqueous solutions. Those properties are outlined below.

1. Aqueous solutions of acids are electrolytes, meaning that they conduct an electrical current. Some acids are strong electrolytes because they ionize completely in water, yielding a great many ions. Other acids are weak electrolytes that exist primarily in a non-ionized form when dissolved in water.
2. Acids have a sour taste. Lemons, vinegar, and sour candies all contain acids.
3. Acids change the color of certain acid-base indicators. Two common indicators are litmus and phenolphthalein. Litmus turns red in the presence of an acid, while phenolphthalein is colorless.
4. Acids react with active metals to yield hydrogen gas. Recall that an activity series is a list of metals in descending order of reactivity. Metals that are above hydrogen in the activity series will replace the hydrogen from an acid in a single-replacement reaction, as shown below. $\mathrm{Zn}(s) + \mathrm{H_2SO_4}(aq) \rightarrow \mathrm{ZnSO_4}(aq) + \mathrm{H_2}(g)$
5. Acids react with bases to produce a salt and water. When equal moles of an acid and a base are combined, the acid is neutralized by the base. Water and an ionic compound called a salt are produced.

Bases

Bases have properties that mostly contrast with those of acids.

1. Aqueous solutions of bases are also electrolytes. Bases can be either strong or weak, just as acids can.
2. Bases often have a bitter taste and are found in foods less frequently than acids. Many bases, like soaps, are slippery to the touch.
3. Bases also change the color of indicators. Litmus turns blue in the presence of a base (Figure below), while phenolphthalein turns pink.
4. Bases do not react with metals in the way that acids do.
5. Bases react with acids to produce a salt and water.

Litmus paper has been treated with the plant dye called litmus. It turns red in the presence of an acid and blue in the presence of a base.

Arrhenius Acids and Bases

Swedish chemist, Svante Arrhenius (1859-1927), was the first to propose a theory to explain the observed behavior of acids and bases. Because of their ability to conduct a current, he knew that both acids and bases contained ions in solution. An Arrhenius acid is a compound which ionizes to yield hydrogen ions (H+) in aqueous solution. An Arrhenius base is a compound which ionizes to yield hydroxide ions (OH) in aqueous solution.

Arrhenius Acids

Acids are molecular compounds with ionizable hydrogen atoms. Only hydrogen atoms that are part of a highly polar covalent bond are ionizable. Hydrogen chloride (HCl) is a gas at room temperature and pressure. The H-Cl bond in hydrogen chloride is a polar bond. The hydrogen atom is electron deficient because of the higher electronegativity of the chlorine atom. Consequently, the hydrogen atom is attracted to the lone pair of electrons in a water molecule when HCl is dissolved in water. The result is that the H-Cl bond breaks, with both bonding electrons remaining with the Cl, forming a chloride ion. The H+ ion attaches to the water molecule, forming a polyatomic ion called the hydronium ion. The hydronium ion (H3O+) can be thought of as a water molecule with an attached hydrogen ion.

Equations showing the ionization of an acid in water are frequently simplified by omitting the water molecule.

HCl(g) → H+(aq) + Cl-

This is merely a simplification of the previous equation, but it is commonly used. Any hydrogen ions in an aqueous solution will be attached to water molecules as hydronium ions.

Not all hydrogen atoms in molecular compounds are ionizable. In methane (CH4), the hydrogen atoms are covalently bonded to carbon in bonds that are only slightly polar. The hydrogen atoms are not capable of ionizing, and methane has no acidic properties. Acetic acid (CH3COOH) (Figure below) belongs to a class of acids called organic acids. There are four hydrogen atoms in the molecule, but only the one hydrogen that is bonded to an oxygen atom is ionizable.

The O-H bond can be ionized to yield the H+ ion and the acetate ion. The other hydrogen atoms in this molecule are not acidic.

The Table below lists some of the more common acids. Recall that the rules for naming acids were covered in the chapter Chemical Nomenclature.

Common Acids
Acid Name Formula
Hydrochloric acid HCl
Nitric acid HNO3
Sulfuric acid H2SO4
Phosphoric acid H3PO4
Acetic acid CH3COOH
Hypochlorous acid HClO

A monoprotic acid is an acid that contains only one ionizable hydrogen. Hydrochloric acid and acetic acid are monoprotic acids. A polyprotic acid is an acid that contains multiple ionizable hydrogens. Most common polyprotic acids are either diprotic (such as H2SO4) or triprotic (such as H3PO4).

Arrhenius Bases

Bases are ionic compounds which yield the hydroxide ion (OH) upon dissociating in water. The Table below lists several of the more common bases.

Common Bases
Base Name Formula
Sodium hydroxide NaOH
Potassium hydroxide KOH
Magnesium hydroxide Mg(OH)2
Calcium hydroxide Ca(OH)2

All of the bases listed in the table are solids at room temperature. Upon dissolving in water, each dissociates into a metal cation and the hydroxide ion.

$\mathrm{NaOH}(s) \mathrm{\overset{H_2O}{\rightarrow} Na^+}(aq) + \mathrm{OH^-}(aq)$

Sodium hydroxide is a very caustic substance also known as lye. Lye is used as a rigorous cleaner and is an ingredient in the manufacture of soaps. Care must be taken with strong bases like sodium hydroxide, as exposure can lead to severe burns (Figure below).

This foot has severe burns due to prolonged contact with a solution of sodium hydroxide, also known as lye.

Sodium belongs to the group of elements called the alkali metals. An alkaline solution is another name for a solution that is basic. All alkali metals react readily with water to produce the metal hydroxide and hydrogen gas. The resulting solutions are basic.

2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)

Bases that consist of an alkali metal cation and the hydroxide anion are all very soluble in water. Compounds of the Group 2 metals (the alkaline earth metals) are also basic. However, these compounds are generally not as soluble in water. Therefore the dissociation reactions for these compounds are shown as equilibrium reactions.

$\mathrm{Mg(OH)_2}(s) \mathrm{\overset{H_2O}{\rightleftharpoons} Mg^{2+}}(aq) + \mathrm{2OH^-}(aq)$

These relatively insoluble hydroxides were some of the compounds discussed in the context of the solubility product constant (Ksp). The solubility of magnesium hydroxide is 0.0084 g per liter of water at 25°C. Because of its low solubility, magnesium hydroxide is not as dangerous as sodium hydroxide. In fact, magnesium hydroxide is the active ingredient in a product called milk of magnesia, which is used as an antacid or a mild laxative (Figure below).

Milk of magnesia is a suspension of magnesium hydroxide, which is nearly insoluble in water. Its low solubility makes it safe to consume.

Brønsted-Lowry Acids and Bases

The Arrhenius definition of acids and bases is somewhat limited. There are some compounds whose properties suggest that they are either acidic or basic, but which do not qualify according to the Arrhenius definition. An example is ammonia (NH3). An aqueous solution of ammonia turns litmus blue, reacts with acids, and displays various other properties that are common for bases. However, it does not contain the hydroxide ion. In 1923, a broader definition of acids and bases was independently proposed by Danish chemist, Johannes Brønsted (1879-1947) and English chemist, Thomas Lowry (1874-1936). A Brønsted-Lowry acid is, a molecule or ion that donates a hydrogen ion in a reaction. A Brønsted-Lowry base is a molecule or ion that accepts a hydrogen ion in a reaction. Because the most common isotope of hydrogen consists of a single proton and a single electron, a hydrogen ion (in which the single electron has been removed) is commonly referred to as a proton. As a result, acids and bases are often called proton donors and proton acceptors, respectively, according to the Brønsted-Lowry definition. All substances that are categorized as acids and bases under the Arrhenius definition are also defined as such under the Brønsted-Lowry definition. The new definition, however, includes some substances that are left out according to the Arrhenius definition.

Brønsted-Lowry Acid-Base Reactions

An acid-base reaction according to the Brønsted-Lowry definition is a transfer of a proton from one molecule or ion to another. When ammonia is dissolved in water, it undergoes the following reversible reaction.

$\underset{\text{base}}{\text{NH}_3}(aq) + \underset{\text{acid}}{\text{H}_2\text{O}}(l) \rightleftharpoons \underset{\text{acid}}{\text{NH}^+_4}(aq) + \underset{\text{base}}{\text{OH}^-}(aq)$

In this reaction, the water molecule is donating a proton to the ammonia molecule. The resulting products are the ammonium ion and the hydroxide ion. The water is acting as a Brønsted-Lowry acid, while the ammonia is acting as a Brønsted-Lowry base. The hydroxide ion that is produced causes the solution to be basic.

We can also consider the reverse reaction in the above equation. In that reaction, the ammonium ion donates a proton to the hydroxide ion. The ammonium ion is a Brønsted-Lowry acid, while the hydroxide ion is a Brønsted-Lowry base. Most Brønsted-Lowry acid-base reactions can be analyzed in this way. There is one acid and one base as reactants, and one acid and one base as products.

In the above reaction, water acted as an acid, which may seem a bit unexpected. Water can also act as a base in a Brønsted-Lowry acid-base reaction, as long as it reacts with a substance that is a better proton donor. Shown below is the reaction of water with the hydrogen sulfate ion.

$\underset{\text{acid}}{\text{HSO}_4^-}(aq) + \underset{\text{base}}{\text{H}_2\text{O}}(l) \rightleftharpoons \underset{\text{acid}}{\text{H}_3\text{O}^+}(aq) + \underset{\text{base}}{\text{SO}^{2-}_4}(aq)$

Water is capable of being either an acid or a base, a characteristic called amphoterism. An amphoteric substance is one that is capable of acting as either an acid or a base by donating or accepting hydrogen ions.

Conjugate Acids and Bases

When a substance that is acting as a Brønsted-Lowry acid donates its proton, it becomes a base in the reverse reaction. In the reaction above, the hydrogen sulfate ion (HSO4) donates a proton to water and becomes a sulfate ion (SO42−). The HSO4 and the SO42− are linked to one another by the presence or absence of the H+ ion. A conjugate acid-base pair is a pair of substances related by the loss or gain of a single hydrogen ion. A conjugate acid is the particle produced when a base accepts a proton. The hydrogen sulfate ion is the conjugate acid of the sulfate ion. A conjugate base is the particle produced when an acid donates a proton. The sulfate ion is the conjugate base of the hydrogen sulfate ion.

A typical Brønsted-Lowry acid-base reaction contains two conjugate acid-base pairs as shown below.

$\mathrm{HNO_2}(aq) + \mathrm{PO^{3-}_4}(aq) \rightleftharpoons \mathrm{NO^-_2}(aq) + \mathrm{HPO^{2-}_4}(aq)$

One conjugate acid-base pair is HNO2 / NO2, while the other pair is HPO42− / PO43−.

Lewis Acids and Bases

Gilbert Lewis (1875-1946) proposed a third theory of acids and bases that is even more general than the Arrhenius and Brønsted-Lowry theories. A Lewis acid is a substance that accepts a pair of electrons to form a covalent bond. A Lewis base is a substance that donates a pair of electrons to form a covalent bond. In a Lewis acid-base reaction, a new covalent bond is formed, but both of the electrons in that bond were originally present on the base. A hydrogen ion, which lacks any electrons, can readily accept a pair of electrons to form a covalent bond. It is an acid under both the Brønsted-Lowry and Lewis definitions. Ammonia has a lone pair on its central nitrogen atom. The reaction between ammonia and the hydrogen ion can be depicted as shown below.

The lone pair on the nitrogen atom provides both electrons for the new covalent N-H bond, so the NH3 is a Lewis base and the H+ is a Lewis acid.

Some reactions that do not qualify as acid-base reactions under the other definitions do so under only the Lewis definition. An example is the reaction of ammonia with boron trifluoride.

Boron trifluoride is the Lewis acid, while ammonia is again the Lewis base. As there is no hydrogen ion involved in this reaction, it qualifies as an acid-base reaction only under the Lewis definition. The Table below summarizes the three acid-base theories.

Acid-Base Definitions
Type Acid Base
Arrhenius H+ ions in solution OH ions in solution
Brønsted-Lowry H+ donor H+ acceptor
Lewis electron-pair acceptor electron-pair donor

Lesson Summary

• Acids and bases have distinct properties. Both are electrolytes and turn indicators a specific color. Acids taste sour, while bases tend to taste bitter. Solutions of acids and bases react with each other to give a neutral solution.
• Arrhenius acids yield H+ ions in aqueous solution. Arrhenius bases yield OH ions in aqueous solution.
• Brønsted-Lowry acids are proton donors, while Brønsted-Lowry bases are proton acceptors. A Brønsted-Lowry acid-base reaction consists of two conjugate acid-base pairs.
• Lewis acids are electron-pair acceptors, while Lewis bases are electron-pair donors.
• The Arrhenius definition of acids and bases is the narrowest definition, while the Lewis definition is the broadest.

Lesson Review Questions

Reviewing Concepts

1. Which statement below is true? Explain.
1. All Arrhenius bases are also Brønsted-Lowry bases.
2. All Brønsted-Lowry bases are also Arrhenius bases.
2. Classify each of the following as an Arrhenius acid, Arrhenius base, or neither.
1. LiOH
2. HClO4
3. H2C2O4
4. CH3COOH
5. Sr(OH)2
6. CH4
3. What does it mean to say that a substance is amphoteric?
4. What must be true about a certain covalent bond in order for a hydrogen atom to be ionizable?
5. In order to be a Brønsted-Lowry base, a molecule or ion must have a lone pair of electrons. Explain why this is true.

Problems

1. Identify each reactant in the following reactions as an acid or a base according to the Brønsted-Lowry theory.
1. $\mathrm{HIO_3}(aq) + \mathrm{H_2O}(l) \rightleftharpoons \mathrm{IO^-_3}(aq) + \mathrm{H_3O^+}(aq)$
2. $\mathrm{F^-}(aq) + \mathrm{HClO}(aq) \rightleftharpoons \mathrm{HF}(aq) + \mathrm{ClO^-}(aq)$
3. $\mathrm{H_2PO^-_4}(aq) + \mathrm{OH^-}(aq) \rightleftharpoons \mathrm{HPO^{2-}_4}(aq) + \mathrm{H_2O}(l)$
4. $\mathrm{CO^{2-}_3}(aq) + \mathrm{H_2O}(l) \rightleftharpoons \mathrm{HCO^-_3}(aq) + \mathrm{OH^-}(aq)$
2. Referring to question 6, identify the conjugate acid-base pairs in each reaction.
3. Write the formula of each acid’s conjugate base.
1. HNO3
2. HSO3
3. H3AsO4
4. HCOOH
4. Write the formula of each base’s conjugate acid.
1. BrO3
2. NH3
3. CH3COO
4. HCO3
5. Explain why the hydrogen phosphate ion (HPO42−) is amphoteric.
6. A positive cation, such as Al3+, undergoes a Lewis acid-base reaction with water. Explain, and identify which species is the Lewis acid and which is the Lewis base.

Points to Consider

The relative acidity or basicity of an aqueous solution is measured using a scale called the pH scale.

• How are the hydrogen and hydroxide ion concentrations related in all aqueous solutions?
• How is the pH of a solution calculated?

Date Created:

Mar 29, 2013

Aug 01, 2014
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