Arrhenius Theory
The Arrhenius Theory, proposed by Svante Arrhenius in 1887, explains the behavior of acids and bases in aqueous solutions. According to this theory, acids are substances that release hydrogen ions (H+) in solution, while bases release hydroxide ions (OH-). This theory laid the foundation for understanding the nature of acids and bases and their interactions in chemical reactions.
6 Key excerpts on "Arrhenius Theory"
eBook - ePub- Arnost Kotyk, Jan Slavik(Authors)
- 2020(Publication Date)
- CRC Press(Publisher)
...IONIZATION OF ACIDS AND BASES Although water is by far the most abundant component of all living systems, its dissociation into oxonium ions is so weak that the pH of a salt solution, either extra- or intracellular, is determined by the presence of components that readily dissociate or readily bind an oxonium ion, i.e., acids and bases, respectively. Throughout modern electrochemistry, three theories of acids and bases came into prominence. The first theory, that of Arrhenius 3 dates back to 1887 when he postulated the universal existence of dissociation of electrolytes in solution, supporting his views by conductometric measurements. He calculated the degree of dissociation α from the ratios of equivalent conductivities at a given and at an infinite dilution. Thus, α = Λ / Λ ∞ It was Arrhenius who defined acids and bases in a simple way, stating that an acid (HA) is characterized by dissociation of hydrogen ions HA ⇌ H + + A − (7a) while a base (BOH) is recognized as a substance dissociating hydroxide ions BOH ⇌ B + + OH − (7b) Although many of his deductions are still valid, particularly with respect to aqueous solutions, he could not foresee the behavior of acids and bases in nonaqueous solvents and the role of interactions between the solute and the solvent. The second theory is the one that is most relevant to our considerations of pH in aqueous solutions and is due to Brönsted 4 who replaced Arrhenius’ definition of acids and bases by stating more generally that an acid is any substance that can dissociate a proton; a base is then any substance that can bind a proton...
eBook - ePub- J Fisher, J.R.P. Arnold, Julie Fisher, John Arnold(Authors)
- 2020(Publication Date)
- Taylor & Francis(Publisher)
...Section N - Aqueous Behavior N1 Lowry—Bronstead acid and base DOI: 10.1201/9780203079522-54 Key Notes Definition In the Lowry—Bronstead theory of acids and bases, an acid is defined as a substance that will give up a hydrogen ion (a proton) and a base is a substance that will accept a proton. This theory superceded that devised by Arrhenius and enabled the concept of acids and bases to extend beyond aqueous solutions. Conjugate acids and bases It is unlikely that protons can exist freely in solution. Consequently if an acid gives up a proton, there must be a base present to accept it. Thus, when an acid and base differ by one proton they are said to be conjugate to each other. Every acid must have its conjugate base and every base its conjugate acid. Ionic strength Dissolved ions, such as those produced by acid—base reactions, cannot be thought of as isolated entities. Rather, the properties of an ion are influenced by the presence of neighboring ions, owing to electrostatic forces between them. Consequently, it follows that the concentration of an ion in solution is not a true reflection of its ability to determine any property of the solution, except at infinite dilution when electrostatic interactions are minimized. To overcome this problem Debye and Hückel defined the term ionic strength, which depends not only on the concentration of all of the ionic species present, but also on the charge carried by each ion. Related topic (C) Water — the biological solvent Definition The terms acid and base were originally defined by Arrhenius as substances that gave rise to hydrogen ions (H +) and hydroxide ions (− OH), respectively, in solution. The combination of an acid and base would lead to neutralization and the formation of neutral water. Of course these definitions were only applicable to aqueous solutions. A more general theory was subsequently developed by Lowry and Bronstead independently...
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
...13 Acids and Bases Chapters 11 and 12 gave you an indication that all acids have certain properties in common and all bases have certain properties in common. The major common property is that acids react with bases (and vice versa) to produce salts. For example, if solutions of HCl (an acid) and KOH (a base) are mixed, the following reaction occurs. Such a reaction gives a solution that no longer has the acidic or basic properties that were evident before mixing, provided the correct volumes and concentrations were used. What then is an acid? What is a base? There are three definitions that have been developed through the years. Each has its own particular usefulness, depending upon the nature of the reactants and the conditions of the reaction. In this chapter we will discuss each of the definitions and their particular usefulness. Our discussion of acids and bases will touch on several other important concepts, including reactions of salts with water, another concentration term specially developed for acid–base solutions, and the importance of acid–base chemistry to physiological and industrial processes. OBJECTIVES After completing this chapter, you will be able to recognize and apply or illustrate the following: Arrhenius, Brønsted–Lowry, and Lewis acids and bases, neutralization, hydrolysis, pH, buffer solution, titration, conjugate acid or base, amphiprotic, indicator, hydronium ion, and hydrated; write a chemical equation for a neutralization reaction between any acid and base; predict whether a solution of a given salt will be acidic, basic, or neutral; calculate the pH of a solution when given: the degree of ionization of a weak acid or base and vice versa, K a or K b of the acid or base and vice versa, the concentration of a solution of a strong acid or base; solve titration problems. ARRHENIUS ACIDS AND BASES There are several chemical theories of acids and bases. The most familiar is that of Arrhenius...
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...Chapter 5 Acids and Bases Abstract This chapter explains the differences between the Brønsted-Lowry and Lewis definitions of acids and bases and gives examples of each. Since chemical reactions involving Lewis acids and bases are covered in more detail in Chapter 10, most of the chapter is dedicated to the applications of the Brønsted-Lowry concepts. This includes the strength of acids and their conjugate bases, the behavior of acids and bases in aqueous solution, the autoionization of water, and the acid ionization constants. The “p” functions, including pH, pOH, p K a, p K b, and p K w, are discussed. The function and uses of buffer solutions are explained along with their design using the Henderson-Hasselbalch equation. Titration procedures are discussed and their relevance to industrial situations is stressed. Keywords Brønsted-Lowry; Lewis acid; Conjugate pairs; Ionization constant; Amphoteric; Coordinate-covalent bond; Coordination complex; Buffer solution; Henderson-Hasselbalch; Titration Outline 5.1 Defining Acids and Bases 5.2 Acids and Bases in Aqueous Solution 5.3 The pH Scale 5.4 Other “p” Functions 5.5 Buffer Solutions 5.6 The Titration Important Terms Study Questions Problems 5.1 Defining Acids and Bases The first modern attempt at defining acids and bases was by a Swedish chemist named Svante Arrhenius in 1887. Arrhenius defined an acid as a material that releases hydrogen ions (H +) when dissolved in water. Similarly, he defined a base as a material that releases hydroxide ions (OH −) when dissolved in water. This definition only held for ionic compounds containing hydrogen or hydroxide ions and did not apply to many acids and bases that we deal with today. Since this early definition of acids and bases was so limited, two more sophisticated and general definitions of acids and bases have since been developed, which are in wide use today. These are known as the Brønsted-Lowry definition and the Lewis definition. In 1923, both J.N...
eBook - ePub- Tony Farine, Mark A. Foss(Authors)
- 2013(Publication Date)
- Routledge(Publisher)
...Since at room temperature only 7% of the ethanoic acid molecules dissociate into ions, the arrow pointing from right to left is emphasised. Hydrogen carbonate (bicarbonate) ion The two previous examples of acids were covalently bonded molecules that dissociated in solution. Now we consider an ion that can act as an acid: We have emphasised the arrows equally to show that, in the body, the equilibrium of the reaction may be tipped to the left or to the right, depending upon the concentration of hydrogen ions. If in low concentration, the reaction equilibrium is tipped to the right and more hydrogen ions are generated. If hydrogen ions are in excess, then the reaction equilibrium will be tipped to the left and hydrogen ions will be removed by combination with carbonate ions. Bases A base is a molecule or ion that accepts a hydrogen ion during a chemical reaction. An older definition of a base is a substance that yields hydroxide ions (OH -) in solution, but in reality the two definitions are saying the same thing. A small number of water molecules dissociate into hydrogen ions and hydroxide ions: Consequently, the addition of any substance that removes hydrogen ions leads to a relative excess of hydroxide ions. Sodium hydroxide Perhaps sodium hydroxide (NaOH) is unfamiliar to you, but you may have heard it referred to by its common name caustic soda. Indeed, you may have even bought products that contain this in order to unblock a drain or clean an oven. The use of sodium hydroxide in such products is related to its ability to dissolve fat; since cell membranes contain a high proportion of fat, sodium hydroxide is potentially dangerous to body tissue, especially if splashed into the eyes. Sodium hydroxide is an ionic compound. Even in the solid state, however, it does not exist as molecules of NaOH but rather as sodium ions (Na +) and hydroxide ions (OH -)...
eBook - ePub- Professor Rob Beynon, J Easterby(Authors)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
...Under physiological conditions, the species on both sides of the equation can co-exist in substantial amounts—compare this with a strong acid such as HCl, which is virtually completely ionized to H + and Cl –. There are other more rigorous definitions of weak acids and bases which were alluded to in Chapter 2, but these need not concern us here. ◊ Nearly all pH buffers are weak acids or bases. Notice that the weak acid can be neutral (acetic acid) or carry a positive (TrisH +) or negative (phosphate 1–) charge. As we develop the theory of buffers, it will become clear that these charges on the buffer species have important consequences. 2. Weak acids and bases resist pH changes A buffer is able to resist changes in pH because it exists in an equilibrium between a form that has a hydrogen ion bound (conjugate acid, protonated) and a form that has lost its hydrogen ion (conjugate base, deprotonated). For the simple example of acetic acid, the equation is: CH 3 COOH ⇌ CH 3 COO − + H + Here, the protonated form is acetic acid, with a net charge of zero, whereas the deprotonated form (acetate) has a charge of −1. The two species are in equilibrium, and this equilibrium, in common with all equilibria, can be displaced by addition of one component. Consider a solution that contains equal amounts of acetic acid and acetate ions (10 mM acetic acid, 10 mM sodium acetate, for example). If we were to add a strong acid, such as HCl, to this solution, the added H + would displace the equilibrium to the left. Binding of H + to CH 3 COO – ‘mops up’ the added protons (Figure 3.1). Electrical neutrality is preserved because every H + that reacts with a CH 3 COO – anion to form the neutral CH 3 COOH leaves behind a chloride (Cl –) anion in its place. Add a strong base, such as sodium hydroxide, and the OH - ion would react with the H + and displace the equilibrium to the right...
