Chemistry
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.
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11 Key excerpts on "Arrhenius Theory"
eBook - PDF- Ambrish Singh(Author)
- 2020(Publication Date)
- IntechOpen(Publisher)
This theory is very limited, out of three theories. According to this theory, the solution medium should be aqueous and acid should produce hydrogen ion (H + ) or base should produce hydroxide ion (OH À ) on dissociation with water. Hence, the substance is regarded as Arrhenius acid or Arrhenius base when it is dissolved in water. For example, HNO 3 is regarded as Arrhenius acid when it is dissolved in aqueous solution. But when it is dissolved in any other solvent like benzene, no dissociation occurs. This is against the Arrhenius Theory. 2. Arrhenius Theory is not applicable on the non-aqueous or gaseous reactions because it explained the acid-base behavior in terms of aqueous solutions. 3. In Arrhenius Theory, salts are produce in the product which are neither acidic nor basic. So, this theory cannot explain the neutralization reaction without the presence of ions. For example, when acetic acid (weak acid) and sodium hydroxide (strong base) reacts, then the resulting solution basic. But this concept is not explained by Arrhenius. 4.Arrhenius Theory is only applicable to those compounds which having formula HA or BOH for acids and bases. There are some acids like AlCl 3 , CuSO 4 , CO 2 , SO 2 which cannot be represented by HA formula, this theory is unable to explain their acidic behavior. Similarly, there are some bases like Na 2 CO 3 , NH 3 , etc. which do not represented by BOH formula, this theory is unable to explain their basic behavior. 9. Bronsted-Lowry theory We have been previously learned an Arrhenius acid-base theory which provided a good start towards the acid-base chemistry but it has certain limitations and problems. After this theory, a Danish chemist, named Johannes Nicolaus Bronsted and British scientist, Thomas Martin Lowry proposed a different definition of acid-base that based on the abilities of compound to either donate or accept the protons. This theory is known as Bronsted-Lowry theory, also called Proton theory of acid and base.
eBook - PDFChemistry for Today
General, Organic, and Biochemistry
- Spencer Seager, Michael Slabaugh, Maren Hansen, , Spencer Seager, Spencer Seager, Michael Slabaugh, Maren Hansen(Authors)
- 2021(Publication Date)
- Cengage Learning EMEA(Publisher)
Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. Acids, Bases, and Salts 293 Concept Summary 9.1 The Arrhenius Theory Learning Objective: Can you write reaction equations that illustrate Arrhenius acid–base behavior? ● Arrhenius acids dissociate in water to provide hydrogen ion (H 1 ). ● Arrhenius bases dissociate in water to provide hydroxide ions (OH 2 ). 9.2 The Brønsted Theory Learning Objective: Can you write reaction equations that illustrate Brønsted acid–base behavior, and identify Brønsted acids and bases from written reaction equations? ● Johannes Bronsted and Thomas Lowry proposed a theory for acids and bases. ● Acids are hydrogen-containing substances capable of donat- ing protons to other substances. ● Bases accept and form covalent bonds with protons. ● When a substance behaves as a Brønsted acid by donating a proton, it becomes a conjugate base. 9.3 Naming Acids Learning Objective: Can you name common acids? ● There are two ways to name acids. ● Water solutions of binary covalent compounds containing hydrogen and a nonmetal are named following the pattern of hydro(stem)ic, where (stem) is the name of the nonmetal bonded to hydrogen. ● Acids in which hydrogen is bonded to polyatomic ions have names based on the name of the polyatomic ion to which the hydrogen is attached. 9.4 The Self-Ionization of Water Learning Objective: Can you do calculations using the concept of the self-ionization of water? ● Water can behave both as a Brønsted acid and a Brønsted base. ● In pure water, water molecules will transfer protons from one water molecule (an acid) to another (the base). 9.5 The pH Concept Learning Objective: Can you do calculations using the pH concept? ● The pH is the negative logarithm of the molar H 3 O 1 concen- tration of a solution. ● Solutions with pH values lower than 7 are acidic. ● Solutions with pH values higher than 7 are basic or alkaline.
eBook - PDFChemistry
An Industry-Based Introduction with CD-ROM
- John Kenkel, Paul B. Kelter, David S. Hage(Authors)
- 2000(Publication Date)
- CRC Press(Publisher)
One final point about this has to do with the viscosities of the concentrated acids. We stated earlier that sulfuric acid is a rather viscous acid. This property actually works to a lab worker’s advantage to a small extent because a highly viscous liquid is less likely to splash. Thus, while the problem of splashing is certainly present with sulfuric acid due to the extreme heat evolution, we need to be equally cautious with other acids as well, especially nitric acid because its viscosity is very low. 12.4 Theories of Acids and Bases 12.4.1 The Arrhenius Theory The view of acids and bases as we have been discussing them thus far is known as the Arrhenius Theory , named after Svante Arrhenius , a Swedish scientist who introduced this theory in 1884. In this theory, an acid is defined as a substance that releases hydrogen ions when dissolved in water. A base is defined as a substance that releases hydroxide ions when dissolved in water. All Arrhenius acids have the symbol for hydrogen first in the formula. See Tables 12.1 and 12.2 again for examples. All Arrhenius bases have hydroxide ions in the formula. All bases listed in Table 12.3 are Arrhenius bases. 12.4.2 The Bronsted-Lowry Theory There are substances other than Arrhenius bases that turn litmus paper blue, feel slippery in water solution, and/or neutralize acids, etc. Thus, Johannes Bronsted , a Danish chemist, and Thomas Lowry , an English chemist, presented a broader definition of acids and bases in 1923. The Bronsted-Lowry H 2 SO 4 2H SO 4 2 → Acids and Bases 319 Theory states that an acid is a proton donor and a base is a proton acceptor. This definition of an acid is similar to that in the Arrhenius Theory (a hydrogen ion is a proton) and, thus, all Arrhenius acids are also Bronsted-Lowry acids and vice versa. However, the Bronsted-Lowry definition of a base is broader. There are substances other than Arrhenius bases that accept hydrogen ions.
eBook - PDF- William H. Brown, Thomas Poon(Authors)
- 2017(Publication Date)
- Wiley(Publisher)
O H O Cl H H H H H C O O C O C – acetic acid hydrochloric acid bicarbonate 2.1 What Are Arrhenius Acids and Bases? The first useful definitions of an acid and a base were put forward by Svante Arrhenius (1859–1927) in 1884; Nobel Prize in Chemistry 1903. According to the original Arrhenius definitions, an acid is a substance that dissolves in water to produce H + ions, and a base is a substance that dissolves in water to produce OH − ions. Today we know that a H + ion does not exist in water because it reacts immediately with an H 2 O molecule to give a hydronium ion, H 3 O + : H aq H O l H O aq Hydronium ion 2 3 Apart from this modification, the Arrhenius definitions of acid and base are still valid and useful today, as long as we are talking about aqueous solutions. However, the Arrhenius concept of acids and bases is so intimately tied to reactions that take place in water that it has no good way to deal with acid–base reactions in nonaqueous solutions. For this reason, we concentrate in this chapter on the Brønsted–Lowry definitions of acids and bases, which are more useful to us in our discussion of reactions of organic compounds. 2.1 Arrhenius acid A substance that dissolves in water to produce H + ions. Arrhenius base A substance that dissolves in water to produce OH − ions. HOW TO 2.1 Use Curved Arrows to Show the Transfer of a Proton from an Acid to a Base We can show the transfer of a proton from an acid to a base by using a symbol called a curved arrow. Curved arrows will be used throughout your study of organic chemistry to describe how reactions proceed. There- fore, it is very important that you become proficient in their use. 1. Write the Lewis structure of each reactant and product, showing all valence electrons on reacting atoms. 2. Use curved arrows to show the change in position of electron pairs during the reaction. The tail of the curved arrow is located at an electron pair.
eBook - PDFChemistry
Structure and Dynamics
- James N. Spencer, George M. Bodner, Lyman H. Rickard(Authors)
- 2011(Publication Date)
- Wiley(Publisher)
Arrhenius argued that bases are compounds that dissociate in water to give OH ions and a positive ion. NaOH is an Arrhenius base because it dissociates in water to give the hydroxide (OH ) and sodium (Na ) ions. An Arrhenius acid therefore can be defined as any substance that ionizes when it dissolves in water to give the hydrogen ion, H . An Arrhenius base is any substance that gives the hydroxide ion, OH , when it dissolves in water. Arrhe- NaOH(s) ¡ H 2 O Na + (aq) + OH - (aq) HCl(g) ¡ H 2 O H + (aq) + Cl - (aq) Zn(s) + 2 HCl(aq) ¡ ZnCl 2 (aq) + H 2 (g) 11.2 THE ARRHENIUS DEFINITION OF ACIDS AND BASES 469 Fig. 11.1 The Arrhenius model assumes that HCl dissociates into H and Cl ions when it dissolves in water. Cl − Cl − Cl − H + H + H + H 2 O nius acids include compounds such as HCl, HCN, and H 2 SO 4 that ionize in water to give the H ion. Arrhenius bases include ionic compounds that contain the OH ion, such as NaOH, KOH, and Ca(OH) 2 . 11.3 The Brønsted–Lowry Definition of Acids and Bases In 1923, Johannes Brønsted and Thomas Lowry independently proposed a more powerful set of definitions of acids and bases. The Brønsted, or Brønsted–Lowry, model is based on the assumption that acids donate H ions to another ion or molecule, which acts as a base. According to this model, HCl doesn’t dissociate in water to form H and Cl ions. Instead, an H ion is transferred from HCl to a water molecule to form an H 3 O ion and a Cl ion. The H 3 O ion is known as the hydronium ion. The Brønsted model of the reac- tion between HCl and water is shown in Figure 11.2. HCl(aq) + H 2 O(l) ¡ H 3 O + (aq) + Cl - (aq) 470 CHAPTER 11 / ACIDS AND BASES Fig. 11.2 The Brønsted model assumes that HCl molecules donate an H ion to water molecules to form H 3 O and Cl ions when HCl dissolves in water. H 3 O + H 3 O + H 3 O + Cl − Cl − Cl − Because it is a proton, an H ion is several orders of magnitude smaller than the smallest atom.
eBook - PDF- Saeed Farrokhpay(Author)
- 2023(Publication Date)
- Arcler Press(Publisher)
Svante Arrhenius (1859–1927), a Swedish chemist, developed the 1 st beneficial theory of acids in 1890: an acidic substance has at least 1 hydrogen atom which may dissociate, or ionize, once mixed with water, generating an anion and a hydrated hydrogen ion: To be considered an “Arrhenius acid,” it should include hydrogen. Many compounds, on the other hand, don’t really comprise hydrogen natively, however when mixed with water, they nonetheless produce hydrogen ions; the hydrogen ions are produced by the water on their own, as a result of the interaction with the material. To provide a very relevant operational concept of an acid, the preceding is provided (Titov, 2016): When a chemical is mixed with water, it produces an oversupply of hydrogen ions, which is known as an acid. When it comes to hydrogen in acids, there will be three critical aspects to remember. Although all Arrhenius acids comprise hydrogen atoms, not over all the atoms of hydrogen in a substance are vulnerable to detachment, hence the –CH 3 hydrogens in acetic acid are classified as “non-acidic.” • One of the most significant aspects of understanding chemistry is the ability to anticipate whether hydrogen atoms in a material would be capable to dissolve at what time. The hydrogen’s that do dissolve may dissolve to varying degrees of intensity. Strong acids, like hydrochloric acid and nitric acid, are efficiently dissolved to a 100% degree in the solution they are in. A limited Basics of Acid-base Chemistry 5 percentage of mostly organic acids, like acetic acid, is dissolved in many solutions; hence, mostly organic acids are weak acids. Fluoric acid and HCN are instances of inorganic acids that are weak (Malhotra, 2018). • Sulfuric acid and phosphoric acid are two popular instances of polyprotic acids that have a large number of ionizable hydrogen atoms are recognized as polyprotic acids. Ampholytes are intermediary types of protons that can both absorb and shed protons, like the HPO 4 2– molecule.
eBook - PDFIntroductory Chemistry
An Active Learning Approach
- Mark Cracolice, Edward Peters, Mark Cracolice(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
This is explained by the reactivity of the hydro- gen ion with metals above hydrogen in the activity series (Table 9.2): Znssd 1 2 H 1 saqd S H 2 sgd 1 Zn 21 saqd A macroscopic property of solutions of bases is that they react with solutions of some metal ions to form solids. Arrhenius Theory explains why. The net ionic equation for the precipitation of magnesium hydroxide is typical: Mg 21 saqd 1 2 OH 2 saqd S MgsOHd 2 ssd The term theory is used in science to describe an explanation for a broad class of related phenomena. Thus, Arrhenius acid–base theory explains many specific acid–base reactions. More acidic More basic The juice of a red cabbage is normally blue-purple. On adding acid, the juice becomes more red. Adding base produces a yellow color. A piece of coral (mostly CaCO 3 ) dissolves in acid to give CO 2 gas. Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas. a b c Figure 17.4 Some characteristic reactions of acids and bases. (a) Red cabbage juice changes color when added to acid or a base solution. (b) Substances containing carbonate ion, such as this piece of coral, react with acids to yield bubbles of carbon dioxide gas as one product of the reaction that occurs. (c) A number of metals react with acid solution to yield hydrogen gas as one product. In this photograph, zinc is reacting with hydrochloric acid. Target Check 17.1 According to the Arrhenius Theory of acids and bases, how do you recognize an acid and a base? 17.3 The Brønsted–Lowry Theory of Acids and Bases Goal 2 Given the equation for a Brønsted–Lowry acid–base reaction, or information from which it can be written, explain how or why it can be so classified. 3 Given the name or formula of a Brønsted–Lowry acid and the name or for- mula of a Brønsted–Lowry base, write the net ionic equation for the reaction between them. Charles D. Winters Copyright 2021 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part.
eBook - ePubInorganic Chemistry for Geochemistry and Environmental Sciences
Fundamentals and Applications
- George W. Luther, III(Authors)
- 2016(Publication Date)
- Wiley(Publisher)
Chapter 7 Acids and Bases 7.1 Introduction Acids and bases have been known since before Roman times, for their ability to transform a chemical substance into other chemical forms. For example, sulfuric acid (, oil of vitriol; vitriolic acid) could be produced by the aqueous oxidation of pyrite, which is a common mineral used for many purposes [1]. dehydrates sugar as well as metal hydrates with sometimes dramatic color changes, and also dissolves metals via redox reactions. It is normally the chemical that is produced and sold in greatest quantity in the world each year. The first base was called lye, which was obtained by leaching ashes with water producing potassium hydroxide solution. Lye is a common name for bases, which are used in a variety of purposes including wood degradation into paper or fibers and soap production. 7.2 Arrhenius and Bronsted–Lowry Definitions Common acids such as hydrochloric acid (HCl) and nitric acid were not formally prepared until about the 16th century, so formal definitions of acids and bases similar to bonding theories are relatively new. In 1884, Svante Arrhenius defined an acid as a chemical species that when dissolved in water produced the hydrogen ion,. Although useful, the definition is limited as it does not encompass a large variety of reactions and only considers water as solvent. In 1923, Brønsted and Lowry defined an acid and base reaction as one that involves a hydrogen ion transfer between two reactants. The acid is the hydrogen ion donor (the Brønsted acid) and the base is the hydrogen ion acceptor (the Brønsted base). This definition applies to all solvents and the gas phase. For example, the reaction between HCl and can occur in water as solvent or in the gas phase (an atmospheric reaction) and results in complete transfer to form the hydronium ion,
eBook - PDF- Young, William Vining, Roberta Day, Beatrice Botch(Authors)
- 2017(Publication Date)
- Cengage Learning EMEA(Publisher)
AlbertSmirnov/iStockphoto.com 17 Acids and Bases Unit Outline 17.1 Introduction to Acids and Bases 17.2 Water and the pH Scale 17.3 Acid and Base Strength 17.4 Estimating the pH of Acid and Base Solutions 17.5 Acid–Base Properties of Salts 17.6 Molecular Structure and Control of Acid–Base Strength In This Unit… We now continue our discussion of chemical equilibria, applying the concepts and techniques developed in Chemical Equilibrium (Unit 16) to the chemistry of acids and bases. In Advanced Acid–Base Equilibria (Unit 18) and Precipitation and Lewis Acid–Base Equilibria (Unit 19) we will continue to study chemical equilibria as it applies to acid–base reactions, buffers, and the chemistry of sparingly soluble compounds. Vasilyev/Shutterstock.com Copyright 2018 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. WCN 02-300 Unit 17 Acids and Bases 528 17.1 Introduction to Acids and Bases 17.1a Acid and Base Definitions We begin our study of acids and bases with the acid and base definitions we introduced in Chemical Reactions and Solution Stoichiometry (Unit 9), the Arrhenius definitions. Arrhenius acid : A substance containing hydrogen that, when dissolved in water, increases the concentration of H 1 ions. Arrhenius base : A substance containing the hydroxide group that, when dissolved in water, increases the concentration of OH 2 ions. The Brønsted–Lowry definition is a broader description of the nature of acids and bases. Brønsted–Lowry acid : A substance that can donate a proton ( H 1 ion). Brønsted–Lowry base : A substance that can accept a proton ( H 1 ion). This definition allows us to define a larger number of compounds as acids or bases and to describe acid–base reactions that take place in solvents other than water (such as ethanol or benzene). Ammonia, NH 3 , for example, is not an Arrhenius base (its formula does not contain a hydroxide group).
eBook - PDF- L. Pataki, E. Zapp, R. Belcher, D Betteridge, L Meites(Authors)
- 2013(Publication Date)
- Pergamon(Publisher)
12 BASIC ANALYTICAL CHEMISTRY In the Arrhenius-Ostwald definition of acids and bases, the two types of compounds are distinguished very rigorously. With the ad-vancement of chemistry the acid-base character of compounds has been found to be a much more general property. The contradictions arising from the old theory can be illustrated by some examples. The diameter of the hydrogen ion, that is of the unsolvated proton, is about 10 ~ 13 cm, being 10 5 times smaller than that of other ions. This small diameter implies such a high mobility according to Stokes' law that the electrolytic conductivity of strong acids would almost be that of the better conducting metals. In reality, however, the mo-bility of the hydrogen ion in solution is several orders of magnitude less than the value corresponding to the free state. The remarkably high hydration energy of the hydrogen ion (837.36 ~k J) also indicates that no free hydrogen ions exist in aqueous solutions, only hydrated forms, for example H+ + H 2 0 ;± H3O+ and H+ + 4H 2 0 ^ H 9 0 4 + . Ammonium salts in liquid ammonia can dissolve certain metals with the evolution of hydrogen; their behaviour in this respect is similar to that of aqueous solutions of strong acids. The effect of the solvent on acid-base character can be illustrated very well by urea, which behaves in anhydrous acetic acid as a strong base, in aqueous solutions as a weak base, but in liquid ammonia as an acid. In certain kinetic investigations, the reactions catalysed by hydro-gen or hydroxide ions were interpreted as being characteristic acid-base catalysed reactions. Subsequent investigations, however, have shown that there are other cations and anions (for example NH 4 + , CH 3 COO~), and even neutral molecules, that can act as acid-base catalysts. As discussed above, free protons do not exist in water. This is generally true for solutions, because protons react with the solvent molecules to form 'onium' ions.
- Peter V. Hobbs(Author)
- 2000(Publication Date)
- Cambridge University Press(Publisher)
For example, HCl reacts with pure liquid ammonia HCl(aq) + NH 3 (1) <=* NHJ (aq) + Cr(aq) (5.6) Since NH 3 has eliminated the acid HCl, we could consider NH 3 as a base. These problems with the Arrhenius Theory led J. Br0nsted and T. Lowry to propose a more general view of acids and bases, in which acids tend to donate protons and bases tend to accept protons. From this view- Strengths of acids and bases; acid-dissociation constant 87 point, in both Reactions (5.5) and (5.6) HC1 acts as an acid, and H 2 O in Reaction (5.5) and NH 3 in Reaction (5.6) act as bases. As indicated by the two-way arrows in Reaction (5.5), H 3 O + (aq) may donate a proton and Cl~(aq) may accept a proton. In this case, Cr(aq) is the base and H 3 O + (aq) the acid. Therefore, we could write HCl(aq) + H 2 O(1) <± H 3 O + (aq) + Cl (aq) acid 1 + base 2 <=* acid 2 + base 1 (5.7) where, HCl(aq) and Cl(aq), which differ only by a proton, are called the conjugate acid-base pair for the forward reaction (indicated by 1), and H 3 O + (aq) and H 2 O(1) are the conjugate acid-base pair for the reverse reaction (indicated by 2). In Reaction (5.6), HC1 and Cl are the con-jugate acid-base pair for the forward reaction, and NHJ and NH 3 are the conjugate acid-base pair for the reverse reaction. 5.4 Strengths of acids and bases; acid-dissociation (or ionization) constant The Br0nsted-Lowry view of acids and bases suggests that the strengths of acids can be compared by measuring their relative tendencies to release a proton to a common base (taken to be water). Thus, if we rep-resent an acid by HA and consider its reaction with water HA(aq) + H 2 O(1) <± H 3 O + (aq) + A(aq) (5.8) we can measure the strength of HA by the equilibrium constant for the forward reaction of (5.8) [H 3 O+(aq)][A-(aq)] *•-PAM] (5 -9) K & is called the acid-dissociation (or ionization) constant for HA.
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