Arrhenius Acid-Base Reactions

This emergency crew is neutralizing an acid spill on the highway by covering it with a basic foam. 

Neutralization reactions keep our bodies in balance and also maintain the health of the world around us. (Photo by Cliff Reiter) 

When an Arrhenius acid is combined with an Arrhenius base, we say that they neutralize each other. By this, we mean that the acid counteracts the properties of the base, and the base counteracts the properties of the acid. For example, a strong acid, such as nitric acid, must be handled with extreme caution, because if it gets on your skin, it could cause severe chemical burns. If you accidentally spilled nitric acid on a laboratory bench, however, you could quickly pour a solution of a weak base, such as sodium hydrogen carbonate, on top of the spill to neutralize the acid and make it safer to wipe. In a similar way, a solution of a weak acid, such as acetic acid, can be poured on a strong base spill to neutralize the base before cleanup. Therefore, reactions between Arrhenius acids and bases are ofren called neutralization reactions. 

Neutralization reactions are important in maintaining the necessary balance of chemicals in your body, and they help keep a similar balance in our oceans and lakes. Neutralization reactions are used in industry to make a wide range of products, including pharmaceuticals, food additives, and fertilizers. Let s look at some of the different forms of Arrhenius acid-base reactions, how they can be visualized, and how to describe them with chemical equations. 

Reactions of Aqueous Strong Arrhenius Acids and Aqueous Strong Arrhenius Bases 

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S. A. Arrhenius defined an acid as any hydrogen-containing species able to release protons and a base as any species able to form hydroxide ions [71]. The aqueous acid-base reaction is the reaction between hydrogen ions and hydroxide ions with water formation. The ions accompanying the hydrogen and hydroxide ions form a salt, so the overall Arrhenius acid-base reaction can be written ... 

As you can see, the water is formed from combining the hydrogen and hydroxide ions. In fact, the net-ionic equation (the equation showing only those chemical substances that are changed during the reaction) is the same for all Arrhenius acid-base reactions ... 

The Lewis definitions of acids and bases provide for a more general view of acid-base reactions than either the Arrhenius or Br0nsted-Lowry pic ture A Lewis acid is an electron pair acceptor A Lewis base is an electron pair donor The Lewis approach incorporates the Br0nsted-Lowry approach as a subcategory m which the atom that accepts the electron pair m the Lewis acid is a proton... 

Thus, the relationship between acid and base is a reciprocal one and an acid-base reaction involves the transfer of a proton. This concept is not restricted to aqueous solutions and it discards Arrhenius prerequisite of ionization. 

J. N. Bronsted and T. M. Lowry independently arrived at definitions of an acid and a base that do not involve water. They recognized that the essential characteristic of an acid-base reaction was the transfer of a hydrogen ion (proton) from one species (the acid) to another (the base). According to these definitions, an acid is a proton donor and a base is a proton acceptor. The proton must be donated to some other species so there is no acid without a base. According to Arrhenius, HC1 is an acid because... 

Acid-base reactions concepts of Arrhenius, Brpnstcd-Lowry, and Lewis coordination complexes, amphoterism... 

Thus, Lewis s definition is a much broader definition that includes coordination compound formation as acid-base reactions, besides Arrhenius and Lowry-Bronsted acids and bases. Examples ... 

The Arrhenius theory explains acid-base reactions as a combination of H (aq) and OH (aq). It provides insight into the heat of neutralization for the reaction between a strong acid and a strong base. (Strong acids and bases dissociate completely into ions in solution.) For example, consider the following reaction. 

The limitations of the Arrhenius theory of acids and bases are overcome by a more general theory, called the Bronsted-Lowry theory. This theory was proposed independently, in 1923, by Johannes Br0nsted, a Danish chemist, and Thomas Lowry, an English chemist. It recognizes an acid-base reaction as a chemical equilibrium, having both a forward reaction and a reverse reaction that involve the transfer of a proton. The Bronsted-Lowry theory defines acids and bases as follows ... 

The Arrhenius theory accounts for the properties of many common acids and bases, but it has important limitations. For one thing, the Arrhenius theory is restricted to aqueous solutions for another, it doesn t account for the basicity of substances like ammonia (NH3) that don t contain OH groups. In 1923, a more general theory of acids and bases was proposed independently by the Danish chemist Johannes Bronsted and the English chemist Thomas Lowry. According to the Bronsted-Lowry theory, an acid is any substance (molecule or ion) that can transfer a proton (H + ion) to another substance, and a base is any substance that can accept a proton. In short, acids are proton donors, bases are proton acceptors, and acid-base reactions are proton-transfer reactions ... 

According to the Arrhenius theory, acids (HA) are substances that dissociate in water to produce H + (aq). Bases (MOH) are substances that dissociate to yield OH aq). The more general Bransted-Lowry theory defines an acid as a proton donor, a base as a proton acceptor, and an acid-base reaction as a proton-transfer reaction. Examples of Bronsted-Lowry acids are HC1, NH4+, and HSO4- examples of Bronsted-Lowry bases are OH-, F-, and NH3. 

Arrhenius acid-base theory - Arrhenius developed the theory of the electrolytic dissociation (1883-1887). According to him, an acid is a substance which delivers hydrogen ions to the solution. A base is a substance which delivers hydroxide ions to the solution. Accordingly, the neutralization reaction of an acid with a base is the formation of water and a salt. It is a so-called symmetrical definition because both, acids and bases must fulfill a constitutional criterion (presence of hydrogen or hydroxide) and a functional criterion (to deliver hydrogen ions or hydroxide ions). The theory could explain all of the known acids at that time and most of the bases, however, it could not explain the alkaline properties of substances like ammonia and it did not include the role of the solvent. -> Sorensen (1909) introduced the -> pH concept. 

There is yet another problem with the Arrhenius theory. It is limited to acid-base reactions in a single solvent, water. Many acid-base reactions take place in other solvents, however. 

The Br0nsted-Lowery theory (usually called the Br0nsted theory), advanced by these workers in 1923, is more comprehensive than the Arrhenius theory. According to this theory, an acid-base reaction is characterized as a reaction in which a proton is transferred from one species (the acid, the proton donor) to another (the base, the proton acceptor). There can be no acid without a base neither exists in isolation because the proton must be transferred to some other species. According to this theory, HC1 is an acid because when it is placed in water, it acts as a proton donor,... 

The Br0nsted-Lowry definition of acids and bases does not replace the Arrhenius definition, but extends it. The Bronsted-Lowry definition of acids and bases requires you to take a closer look at the reactants and products of an acid-base reaction. In this case, acids and bases are not easily defined as having hydronium and hydroxide ions. Instead, you are asked to look and see which substance has lost a proton and which has gained the very same proton that was lost. 

Note that this reaction is not considered an acid-base reaction according to the Arrhenius concept. 

The Lewis bonding model with its electron pairs can be used to define a more general kind of acid-base behavior of which the Arrhenius and Bronsted-Lowry definitions are special cases. A Lewis base is any species that donates lone-pair electrons, and a Lewis acid is any species that accepts such electron pairs. The Arrhenius acids and bases considered so far fit this description (with the Lewis acid, H, acting as an acceptor toward various Lewis bases such as NH3 and OH , the electron pair donors). Other reactions that do not involve hydrogen ions can still be considered Lewis acid-base reactions. An example is the reaction between electron-deficient BF3 and electron-rich NH3 ... 

Although oxoacids and hydroxides are Arrhenius acids and bases (they release or OYi (aq) into aqueous solution), acid and base anhydrides do not fall into this classification because they contain neither nor OH. Acid anhydrides are acids in the Lewis sense, (they accept electron pairs), and base anhydrides are bases in the Lewis sense, (their ions donate electron pairs). The reaction between an acid anhydride and a base anhydride is then a Lewis acid-base reaction. An example of such a reaction is... 

Note that these reactions are not redox reactions (oxidation numbers do not change). They are clearly not dissolution or precipitation reactions, nor are they acid-base reactions in the Arrhenius sense. But they are usefully classified as acid-base reactions in the Lewis sense. 

If ammonia had been allowed to react with hydrogen chloride in the aqueous phase, then it would have fallen within the Arrhenius definition of an acid/base reaction. Notice that it is the water molecule that is central to the Arrhenius definition, and this is one of the main reasons why it is of little use in conventional organic chemistry. Similarly, because the Bronsted-Lowry definition does not depend on water, but only upon the presence of two compounds of complementary properties, it is this definition that is of more use to organic chemists. 

In 1884, Svante Arrhenius (1859-1927) presented his theory of electrolytic dissociation, which resulted in the Arrhenius theory of acid-base reactions. In his view. 

Arrhenius and Bronsted-Lowry acid-base neutralization reactions all have one thing in common. They involve the reaction of an acid with a base to form a salt that contains the cation characteristic of the base and the anion characteristic of the acid. Water is also usually formed. This is indicated in the formula unit equation. The general form of the net ionic equation, however, is different for different acid-base reactions. The net ionic equations depend on the solubility and extent of ionization or dissociation of each reactant and product. 

To this point we have examined acid-base reactions in which stoichiometric amounts of Arrhenius acids and bases were mixed. Those reactions form normal salts. As the name implies, normal salts contain no ionizable H atoms or OH groups. The complete neutralization of phosphoric acid, H3PO4, with sodium hydroxide, NaOH, produces the normal salt, Na3 04. The equation for this complete neutralization is... 

Many organic and biological reactions are acid-base reactions that do not lit within the Arrhenius or Bronsted-Lowry theories. Experienced chemists find the Lewis theory to be very usefol because so many other chemical reactions are covered by it. The less experienced sometimes find the theory less useful, but as their knowledge expands so does its utility. 

Although the Arrhenius definitions of acid, base, and acid-base reaction are very useful, an alternate set of definitions is also commonly employed. In this system, a Bronsted-Lowry acid is a proton (H+) donor, a Bronsted-Lowry base is a proton acceptor, and a Bronsted-Lowry acid-base reaction is a proton transfer. Table 5.7 summarizes the definitions of acid and base in the Arrhenius and Bronsted-Lowry systems. 

These reactions are very similar, but only the first reaction would be considered an acid-base reaction in the Arrhenius system. In each of the reactions, an H" " is transferred from one reactant to another, but only the first is a reaction between an Arrhenius acid and an Arrhenius base. In the first reaction, an is transferred from the Arrhenius weak acid acetic acid, HC2H302( ), to the Arrhenius weak base ammonia, NH3(<3 ). 

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