Acids and Bases

Acids and bases have been known for a long time. When Robert Boyle characterized them in 1680, he noted that acids dissolve many substances, change the color of certain natural dyes (for example, they change litmus from blue to red), and lose these characteristic properties after coming into contact with alkali (bases). In the eighteenth century, it was recognized that acids have a sour taste, react with limestone to liberate a gaseous substance (now known to be CO2 ), and interact with alkali to form neutral substances. In 1815, Humphry Davy contributed greatly to the development of the modern acid-base concept by demonstrating that hydrogen is the essential constituent of acids. Around that same time, Joseph Louis Gay-Lussac concluded that acids are substances that can neutralize bases and that these two classes of substances can be defined only in terms of each other. The significance of hydrogen was reemphasized in 1884 when Svante Arrhenius defined an acid as a compound that dissolves in water to yield hydrogen cations (now recognized to be hydronium ions) and a base as a compound that dissolves in water to yield hydroxide anions.

Acids and bases are common solutions that exist everywhere. Almost every liquid that we encounter in our daily lives consists of acidic and basic properties, with the exception of water. Acids and bases in aqueous solutions will conduct electricity because they contain dissolved ions. Therefore, acids and bases are electrolytes. Strong acids and bases will be strong electrolytes. Weak acids and bases will be weak electrolytes. This affects the amount of conductivity.They have completely different properties and are able to neutralize to form H2O, which will be discussed later in a subsection. Acids and bases can be defined by their physical and chemical observations as follow :

Here are three definitions of acids and bases :

The Arrhenius Definition of Acids and Bases

The Arrhenius definition of acid-base reactions is a development of the "hydrogen theory of acids". Acids are substances that ionized in an aqueous solution to produce hydrogen ions while bases produces hydroxide ions. The Arrhenius definition can only describe acids and bases in an aqueous environment.

Brønsted Acids and Bases

a Brønsted acid as a substance capable of donating a proton, and a Brønsted base as a substance that can accept a proton. An extension of the Brønsted definition of acids and bases is the concept of the conjugate acid-base pair, which can be defined as an acid and its conjugate base or a base and its conjugate acid. The conjugate base of a Brønsted acid is the species that remains when one proton has been removed from the acid. Conversely, a conjugate acid results from the addition of a proton to a Brønsted base. The Brønsted definition also enables us to classify ammonia as a base because of its ability to accept a proton. In this case, NH4+ is the conjugate acid of the base NH4+, and the hydroxide ion OH- is the conjugate base of the acid H2O. Take a look the following reaction :

The Acid-Base Properties of Water

Water, as we know, is a unique solvent. One of its special properties is its ability to act either as an acid or as a base. It can be act as acid if reacts with bases and also can be act as base if react with acids. Water is a very weak electrolyte and therefore a poor conductor of electricity, but it does undergo ionization to a small extent. This reaction is sometimes called the autoionization of water. The equilibrium constant for the ionization of water is called the ion-product constant for water Kw (1x10-14 ).

Like water, many molecules and ions may either gain or lose a proton under the appropriate conditions. Such species are said to be amphiprotic. Another term used to describe such species is amphoteric, which is a more general term for a species that may act either as an acid or a base by any definition (not just the Brønsted-Lowry one). Consider for example the bicarbonate ion, which may either donate or accept a proton as shown here :

pH—A Measure of Acidity

Because the concentrations of H+ and OH- ions in aqueous solutions are frequently very small numbers and therefore inconvenient to work with, Soren Sorensen in 1909 proposed a more practical measure called pH. The pH of a solution is defined as the negative logarithm of the hydrogen ion concentration (in mol/L).

pH = - log [H3O+] or pH = -log [H+]

Because pH is simply a way to express hydrogen ion concentration, acidic and basic solutions at 25°C can be distinguished by their pH values, as follows:

a. Acidic solutions: [H+], <1.0 x 10-7M, pH < 7.00

b. Basic solutions: [H+] , > 1.0 x 10-7M, pH >7.00

c. Neutral solutions: [H+] = 1.0 x 10-7M, pH = 7.00

A pOH scale analogous to the pH scale can be devised using the negative logarithm of the hydroxide ion concentration of a solution. Considering ion -product constant at 250C we can get : pH + pOH = 14.

Strength of Acids and Bases

Strong acids are strong electrolytes that, for practical purposes, are assumed to ionize completely in water to give H+ ions or H3O+ ions . Most of the strong acids are inorganic acids: hydrochloric acid (HCl), nitric acid (HNO3), perchloric acid (HClO4), and sulfuric acid (H2SO4). While weak acids are ionize only to a limited extent in water. At equilibrium, aqueous solutions of weak acids contain a mixture of non-ionized acid molecules, H3O+ ions, and the conjugate base. Examples of weak acids are hydrofluoric acid (HF), acetic acid (CH3COOH), and the ammonium ion (NH4+).

Like strong acids, strong bases are strong electrolytes that ionize completely in water to give hydroxide ions. Hydroxides of alkali metals and certain alkaline earth metals are strong bases. Weak bases, like weak acids, are weak electrolytes. Ammonia is a weak base. It ionizes to a very limited extent in water. Note that, unlike acids, NH3 does not donate a proton to water. Rather, NH3 behaves as a base by accepting a proton from water to form NH4+ and OH- ions.

Conjugate acid-base pairs have the following properties :

If an acid is strong, its conjugate base has no measurable strength. Thus, the Cl- ion, which is the conjugate base of the strong acid HCl, is an extremely weak base.
H3O+is the strongest acid that can exist in aqueous solution. Stronger scids than H3O+ react with water to produce H3O+ and their conjugate bases. Thus, HCl, which is a stronger acid than H3O+, reacts with water completely to form H3O+ and Cl-. While weaker acids will react with water to much smaller extent, producing H3O+ and their conjugate bases.
The OH- ion is the strongest base that can exist in aqueous solution. Bases stronger than OH- react with water to produce OH- and their conjugate acids.

Weak Acids and Acid Ionization Constants

Ka, the acid ionization constant, is the equilibrium constant for the ionization of an acid. At a given temperature, the strength of the acid HA is measured quantitatively by the magnitude of Ka. The larger Ka, the stronger the acid—that is, the greater the concentration of H+ ions at equilibrium due to its ionization. The equilibrium expression for this ionization is :

Another measure of the strength of an acid is its percent ionization. The stronger the acid, the greater the percent ionization. Percent ionization is ionized acid concentration at equilibrium over initial concentration of acid.

Weak Bases and Base Ionization Constants

The ionization of weak bases is treated in the same way as the ionization of weak acids. Thus, the base ionization constant (Kb), which is the equilibrium constant for the ionization reaction.

The Relationship Between the Ionization Constants of Acids and Their Conjugate Bases

When two reactions are added to give a third reaction, the equilibrium constant for the third reaction is the product of the equilibrium constants for the two added reactions. Thus, for any conjugate acid-base pair it is always true that Ka.Kb = Kw.

Diprotic and Polyprotic Acids

Diprotic acid is an acid that has two ionized hydrogen atoms, while polyprotic acid has more than two (poly) ionized hydrogen atoms. E.g for diprotic and polyprotic consecutively Oxalic acid and Phosphoric acid. Because it has more than one hydrogen atom, both diprotic and polyprotic acids have more than one ionization constant, Ka1, Ka2,... so on depends on the number of hydrogen atoms produced. Ionization constant decreased markedly for the second and the third. The trend happened because it is easier to remove H+ ion from neutral molecule than remove H+ ion from a negative ion derived from the molecule.

Molecular Structure and the Strength of Acids

The strength of an acid depends on a number of factors, such as the properties of the solvent, the temperature, and, of course, the molecular structure of the acid. When we compare the strengths of two acids, we can eliminate some variables by considering their properties in the same solvent and at the same temperature and concentration. Then we can focus on the structure of the acids. Two factors influence the ionization of acid are the strength of H-X bond and the polarity of H-X bond. The stronger H-X bond, the ore difficult for HX molecule to break up hence the weaker of acid, while the higher polarity it tends the accumulation of positive and negative charges on H atom and X atom, means that the stronger acid. Now lets discuss about three kinds of acid :

Hydrohalic Acids

Based on the electronegativity in halogen group and the table of bond enthalpy we get that the series of acid from the strongest to the weakest in hydrohalic acid is HF<HCl<HBr<HI. It indicates that bond enthalpy is the predominant factor in determining the strength of binary acid.

Oxoacids

It contains oxygen, hydrogen and other element, e.g phosphoric acid, sulfuric acid etc. Remember the more electronegative element the higher oxidation state that it has, so it will make covalent bond between the element and the oxygen atom and the O-H bond more polar. The consequent is the tendency of the molecule increases to release H+ion. To compare their strengths, it is convenient to divide the oxoacids into two groups:

a. Oxoacids Having Different Central Atoms That Are from the Same Group of the Periodic Table and That Have the Same Oxidation Number. Why does HClO3 is stronger than HBrO3 ? Even though both of them have the same of oxidation number, since Cl is more electronegative than Br so the O-H bond in HClO3 is more polar than in HBrO3 .

b. Oxoacids Having the Same Central Atom but Different Numbers of Attached Groups. Because the number of oxygen atom attached on Cl is larger, the ability of Cl to draw the electron away from Oh group increases, so the series are : HClO4 > HClO3 > HClO2 > HClO.

Carboxylic Acids

The strength of carboxylic acids depends on the nature of the R group. Consider, for example, acetic acid and chloroacetic acid. The presence of the electronegative Cl atom in chloroacetic acid shifts electron density toward the R group, thereby making the OOH bond more polar. Consequently, there is a greater tendency for the acid to ionize.

The conjugate base of the carboxylic acid, called the carboxylate anion (RCOO-), can exhibit resonance. In the language of molecular orbital theory, we attribute the stability of the anion to its ability to spread or delocalize the electron density over several atoms. The greater the extent of electron delocalization, the more stable the anion and the greater the tendency for the acid to undergo ionization. Thus, benzoic acid is a stronger acid than acetic acid because the benzene ring facilitates electron delocalization, so that the benzoate anion (C6H5COO-) is more stable than the acetate anion (CH3COO-).

Acid-Base Properties of Salts

A salt is an ionic compound formed by the reaction between an acid and a base. Salts are strong electrolytes that completely dissociate into ions in water. The term salt hydrolysis describes the reaction of an anion or a cation of a salt, or both, with water. Salt hydrolysis usually affects the pH of a solution.

A. Salt that produce neutral solutions

The salts that contain an alkali metal ion or alkaline earth metal ion (except Be2+) and the conjugate base of a strong acid (for example, Cl2, Br2, No3-) do not undergo hydrolysis to an appreciable extent, and their solutions are assumed to be neutral. E.g NaNO3 the hydrated sodium ion neither donates nor accept proton and the conjugate based is derived from strong acid (nitric acid) an dit has no affinity for proton.

B. Salt that produce basic solution

The solution of a salt derived from a strong base and a weak acid is basic. Because the conjugate base derived from weak acid so it has affinity for H+ion. The salt (e.g sodium acetic) hydrolysis reaction is given by : CH3COO-(aq) + H2O(l) --> CH3COOH(aq) + OH-(aq)

Because this reaction produces OH- ions, the sodium acetate solution will be basic.

C. Salt that produce acidic solution

The solution of a salt derived from strong acid and weak base is acidic. Because the conjugate acid derived from weak base so it donates H+ ion. The salt (e.g ammonium chloride) hydrolysis reaction is given by : NH4+(aq) + H2O(l) à NH3(aq) + H3O+(aq)

Because this reaction produce H+ ion, so ammonium chloride will be acidic. Beside that there are the other cases in determining the pH of salt solution, take a look on the following table.

Acid-Base Properties of Oxides and Hydroxides

Oxides can be classified as acidic, basic, or amphoteric. The following figure will tell you the acidic properties of the oxides. The reaction between CO2 and H2O explains why when pure water is exposed to air (which contains CO2 ) it gradually reaches a pH of about 5.5.

Lewis Acids and Bases

In 1932 the American chemist G. N. Lewis formulated such a definition. He defined what we now call a Lewis base as a substance that can donate a pair of electrons. A Lewis acid is a substance that can accept a pair of electrons. E.g in the protonation of ammonia, NH3 acts as a Lewis base because it donates a pair of electrons to the proton H+ , which acts as a Lewis acid by accepting the pair of electrons.

The significance of the Lewis concept is that it is more general than other definitions. Lewis acid-base reactions include many reactions that do not involve Brønsted acids. E.g the reaction between BF3 and NH3 . The vacant, un-hybridized 2pz orbital in B atom accepts the pair of electrons from NH3 . So BF3 functions as an acid according to the Lewis definition, even though it does not have an ionization ability.

Published by Galuh Catur Wisnu Prabowo, on Friday, August 9, 2019.

Reference : Chang, R. (2010). Chemistry 10th edition (Chapter 15 page 658-697). New York: McGraw-Hill.

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