Acid — Base Reactions

These reactions occur by proton transfer. The protonated or nonprotonated species or perhaps the intermediate species react with the other components of the reaction. According to Bronsted, the following reaction scheme occurs [Pg.154]

Bronsted + Conjugated base Conjugated acid + Acid [Pg.154]

Unlike the previous case, this reaction does not involve electrons, and, therefore, there is no repulsion. The molecules get polarized and have a high degree of solvation in a polar solvent. [Pg.154]

The rate or rate constant follows the theory of Bronsted-Lowry and is a function of the ions, in particular of hydronium ion, which is a combination of the proton with water. It also depends on the concentrations of conjugate acids. Its expression is [Pg.154]

This constant depends on the temperature and mainly on the pH of the solution. When weak acids and bases are involved, we can simply have. [Pg.154]

Acids and bases—dissociation of strong and weak acids. [Pg.128]

Many neutral and all anionic clusters are Lewis bases. Dependent upon the nucleophilic character of the metal atoms and ligands, addition of a proton may occur either to the metal or to the ligand. Usually, M —H bonds are formed first. There are known cases of addition of two protons to neutral electron-rich clusters, and the formation of anionic clusters possessing charges ranging from -1 to -6, for example, [Ni3sPt6(CO)4s] -. [Pg.187]

Reactions with acids sometimes lead to protonation of ligands. Bridging ligands such as H2-, and ii -CO undergo protonation with particular ease [equations (3.78) and (3.79)]. [Pg.187]

In carbyne osmium compounds such as [Os3(/i-COMe)(/z-H)(CO)io] nucleophilic attack on the carbyne carbon atom also takes place. By carrying out sequential reactions with nucleophiles and electrophiles, it is possible to break the C—O bond (Table 3.14, footnote reference A). The reaction furnishes an isolable carbene complex, [Os3(M-CHOMe)(/i-H)(CO)io] [Pg.189]

However, a different reaction mechanism operates for another nucleophile-electrophile pair, Ph and see scheme (3.85). [Pg.189]

Substitution of CO ligands in clusters is most commonly realized in the same way as in the case of mononuclear metal carbonyls. Substitution may be induced by one of the following, most frequently utilized methods thermal, electrochemical, chemical (reactions with N-oxide of trimethylamine or Bu PO), photochemical, catalysis by radicals, catalysis by transition metal compounds, etc. [Pg.190]

According to the general acid-base concepts of Bronsted and Lewis, metal cations are generally regarded as acids. Therefore, transition metal cations or coordina-tively imsaturated compoimds can undergo addition of neutral or anionic nucleophiles to give cationic (Eq. 2-17), anionic (Eq. 2-18), and ir-acceptor complexes (Eq. 2-19). [Pg.21]

Another example of Lewis acid behavior is shown in Equation 2-20, in which an iridium complex takes up a CO hgand to form a dicarbonyl complex. [Pg.22]

In the reverse of dissociation, 16-electron species can add a ligand to give 18-elec-tron complexes [19] [Pg.22]

The Bronsted theory states that the acid/base character of a compound depends on its reaction partner and is therefore not an absolute. An indication that transition metal compounds can act as bases is provided by the long-known protonation reactions of transition metal complexes, generally of low oxidation state. An example is cobalt carbonyl hydride, the true catalyst in many carbonylation reactions [Pg.22]

Metal basicity is also exhibited by phosphine and phosphite complexes of nickel(0), which can be protonated by acids of various strengths [Pg.22]

A useful definition of acids and bases is that independently introduced by Johannes Bronsted (1879-1947) and Thomas Lowry (1874-1936) in 1923. In the Bronsted-Lowry definition, acids are proton donors, and bases are proton acceptors. Note that these definitions are interrelated. Defining a base as a proton acceptor means an acid must be available to provide the proton. For example, in reaction 6.7 acetic acid, CH3COOH, donates a proton to ammonia, NH3, which serves as the base. [Pg.140]

When an acid and a base react, the products are a new acid and base. For example, the acetate ion, C1T3COO-, in reaction 6.7 is a base that reacts with the acidic ammonium ion, N1T45 , to produce acetic acid and ammonia. We call the acetate ion the conjugate base of acetic acid, and the ammonium ion is the conjugate acid of ammonia. [Pg.140]

The equilibrium constant for a reaction in which an acid donates a proton to the solvent (Vi). [Pg.140]

Strong and Weak Acids The reaction of an acid with its solvent (typically water) is called an acid dissociation reaction. Acids are divided into two categories based on the ease with which they can donate protons to the solvent. Strong acids, such as Fid, almost completely transfer their protons to the solvent molecules. [Pg.140]

Weak acids, of which aqueous acetic acid is one example, cannot completely donate their acidic protons to the solvent. Instead, most of the acid remains undissociated, with only a small fraction present as the conjugate base. [Pg.140]

Acids and bases are extremely common, as are the reactions between acids and bases. The driving force is often the hydronium ion reacting with the hydroxide ion to form water. The chapter on Equilibrium describes the equilibrium reactions of acids and bases, as well as some information concerning acid—base titration. After you finish this section, you may want to review the acid-base part of the Equilibrium chapter. [Pg.76]

At the macroscopic level, acids taste sour, may be damaging to the skin, and react with bases to yield salts. Bases taste bitter, feel slippery, and react with acids to form salts. [Pg.76]

At the microscopic level, acids are defined as proton (H ) donors (Bronsted-Lowry theory) or electron-pair acceptors (Lewis theory). Bases are defined as proton (H+) acceptors (Bronsted-Lowry theory) or electron-pair donors (Lewis theory). Consider the gas-phase reaction between hydrogen chloride and ammonia [Pg.76]

HC1 is the acid, because it is donating an H+ and the H+ will accept an electron pair from ammonia. Ammonia is the base, accepting the H+ and furnishing an electron pair that the H+ will bond with via coordinate covalent bonding. Coordinate covalent bonds are covalent bonds in which one of the atoms furnishes both of the electrons for the bond. After the bond is formed, it is identical to a covalent bond formed by donation of one electron by both of the bonding atoms. [Pg.76]

Acids and bases may be strong, dissociating completely, or weak, partially dissociating and forming an equilibrium system. (See Chapter 15 for the details on weak acids and bases.) Strong acids include [Pg.76]

The Br0nsted-Lowry concept of acids and bases will be discussed in detail in Chapter 14. [Pg.154]

Earlier in this chapter we considered Arrhenius s concept of acids and bases An acid is a substance that produces ions when dissolved in water, and a base is a substance that produces OH ions. Although these ideas are fundamentally correct, it is convenient to have a more general definition of a base, which includes substances that do not contain OEI ions. Such a definition was provided by Johannes N. Bronsted (1879-1947) and Thomas M. Lowry (1874-1936), who defined acids and bases as follows [Pg.154]

How do we know when to expect an acid-base reaction One of the most difficult tasks for someone inexperienced in chemistry is to predict what reaction might occur when two solutions are mixed. With precipitation reactions, we found that the best way to deal with this problem is to focus on the species actually present in the mixed solution. This idea also applies to acid-base reactions. For example, when an aqueous solution of hydrogen chloride (HCI) is mixed with an aqueous solution of sodium hydroxide (NaOH), the combined solution contains the ions H, Cl, Na, and OH . The separated ions are present because HCI is a strong acid and NaOH is a strong base. How can we predict what reaction occurs, if any First, will NaCI precipitate From Table 4.1 we can see that NaCI is soluble in water and thus will not precipitate. Therefore, the Na and Cl ions are spectator ions. On the other hand, because water is a nonelectrolyte, large quantities of H and OH ions cannot coexist in solution. They react to form H2O molecules [Pg.154]

This is the net ionic equation for the reaction that occurs when aqueous solutions of HCI and NaOH are mixed. [Pg.154]

Copyright 2010 Cengage Learning, Inc. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. [Pg.154]

This reaction illustrates a very important general principle The hydroxide ion is such a strong base that for purposes of stoichiometric calculations it can be assumed to react completely with any weak acid that we will encounter. Of course, OH ions also react completely with the H+ ions in solutions of strong acids. [Pg.163]

The pKa-values can be used to predict if an acid-base reaction can take place. An acid will donate a proton to the conjugate base of any acid [Pg.16]

How do we recognize acid-base reactions One of the most difficult tasks for someone inexperienced in chemistry is to predict which reaction might occur when two solutions are mixed. With precipitation reactions we found that [Pg.110]

We will now deal with the stoichiometry of acid-base reactions in aqueous solutions. The procedure is fundamentally the same as that used previously. [Pg.111]

Two especially important categories of aqueous solutions are acids and bases. Examples of them are easy to find in our everyday lives as well as in the chemical industry. For our present purposes, we will define an acid as any substance that dissolves in water to produce H (or H3O+) ions and a base as any substance that dissolves in water to produce OH ions. Table 3.2 lists some common acids and bases. Like other solutes, acids and bases can be either strong or weak electrolytes. [Pg.95]

HNO3 Nitric acid NaOH Sodium hydroxide [Pg.96]

Note All common strong acids and bases are shown, but only representative examples of weak acids and bases are listed. [Pg.96]

Unlike ions, hydroxide ions are not represented as having combined with water from the solvent, so water doesn t appear In the equation for a strong base. [Pg.96]

We will revisit the concepts of acids and bases in Chapter 12 and provide additional detail at that time. [Pg.96]

Reaction between a metal ion (electron-pair acceptor, Lewis acid) and a Lewis base (electron-pair donor). For example, complex formation (see Chapter 3) [Pg.15]

This is called the water dissociation constant, Kw (its value at 25°C is approximately 1 x 10-14). [Pg.15]

Example 2.2 Calculate the equilibrium concentration of [OH ] in a water sample where [Pg.15]

Example 2.3 Calculate the solubility of CaS04(S) assuming that the dissolution equilibrium is the only equilibrium present in this water/salt system. Here, KSp = 10-5 9. [Pg.15]

Because the concentration of the cation and the anion is the same (let us call it ), then [Pg.15]

Vinegar contains a weak acid called acetic acid (HC HjOj). Some drain openers contain lye, which is sodium hydroxide (NaOH) and is a strong base. If we mix the two together, the following reaction takes place [Pg.78]

This reaction is different from the reaction that formed CaCOj because there are no soUds in this reaction. The label aq reminds us that these substances are dissolved in water. Substances that are dissolved in water can be molecular or ionic, but are not sohds. By definition, solids are insoluble in water. [Pg.78]

One of the most important types of chemical reactions is the acid-base reaction. However, the definition of which species constitute acids or bases has evolved over the years as the breadth of known chemical reactions has continued to proliferate. For this reason, it is necessary to first introduce the more common historical definitions of acids and bases so that we may better understand how they each fit into the lexicon of chemical reactivity. Just as there were several complimentary models to facilitate our understanding of chemical bonding, so too there are numerous definitions of what it means to be an acid or a base. Which of these definitions we choose will depend on the complexity of the specific acid-base interaction at hand. Ultimately, however, every acid-base reaction entails a change in the way that the valence electrons are arranged in the atomic or molecular orbitals of the participating species. Therefore, the most modem definition of acid-base chemistry builds upon the MO concepts developed in previous chapters and provides the context for a natural continuation of that discussion. [Pg.455]

The traditional Arrhenius definition of acids and bases derives from the early experiments of Arrhenius and Ostwald on the theory of electrolytic dissociation, specifically as it applies to the autoionization of water given by Equation (I4.1). At 298 K, = 1.0 X 10 and the equilibrium concentrations of H3O+ and OH are identical. [Pg.455]

According to the Arrhenius definition, an acid is any species that increases the concentration of H or removes OH and a base is any species that increases the concentration of OH or removes H+. In reality, however, H+ cannot exist in isolation in aqueous solution, as the free energy for the solvation reaction in Equation (14.2) is extremely exergonic [Pg.455]

With this caveat, we will use the terms H and H30 (the hydronium ion) interchangeably. Using the Arrhenius definition, HCI (g) is an acid because it generates H3O+ (aq) ions when it is dissolved in aqueous solution, according to Equation (14.3). Boric acid is also an Arrhenius acid, as it removes OH ions, as shown in Equation (14.4). On the other hand, NaOH is an Arrhenius base, as it increases the concentration of the hydroxide ion in aqueous solution. [Pg.455]

Arrhenius acids and bases represent a subset of a more generalized model of acid-base theory known as solvent theory. In solvent theory, an acid is defined as any substance that increases the concentration of the cationic species that results from autoionization of the solvent, whereas a base increases the concentration of the anionic species from autoionization. Table 14.1 lists the ion products for some of the more common solvents that undergo autoionization. [Pg.455]

The species that give these solutions their characteristic properties are called acids and bases. In this chapter, we use the definitions first proposed by Svante Arrhenius more than a century ago. [Pg.81]

Acidic and basic household solutions. Many common household items, including vinegar, orange juice, and cola drinks, are acidic. In contrast baking soda and most detergents and cleaning agents are basic. [Pg.81]

An acid is a species that produces H+ ions in water solution. [Pg.81]

We will consider more general definitions of acids and bases in Chapter 13 (Bransted-Lowry) and Chapter 15 (Lewis). [Pg.81]

There are two types of acids, strong and weak, which differ in the extent of their ionization in water. Strong acids ionize completely, forming H+ ions and anions. A typical strong acid is HCl. It undergoes the following reaction on addition to water [Pg.97]

Two other important classes of reactions that occur in aqueous solution are acid-base reactions and gas-evolution reactions. In an acid-base reaction (also called a neutralization reaction), an acid reacts with a base and the two neutralize each other, producing water (or in some cases a weak electrolyte). In a gas-evolution reaction, a gas forms, resulting in bubbling. In both cases, as in precipitation reactions, the reactions occur when the anion from one reactant combines with the cation of the other. Many gas-evolution reactions are also acid-base reactions. [Pg.168]

Recall from Chapter 3 that an add forms ions in solution, and we saw earlier that a base is a substance that produces OH ions in solution. More formally [Pg.168]

These deimitions of acids and bases, called the Arrhenius definitions, are named after Swedish chemist Svante Arrhenius (1859-1927). In Chapter 15, we will learn more general definitions of acids and bases, but these definitions are sufficient to describe neutralization reactions. [Pg.169]

According to the Arrhenius definition, HCl is an acid because it produces ions in solution [Pg.169]

Chemists use W aq) and H-iO iaq) interchangeably to mean the same thing—a hydronium ion. The chemical equation for the ionization of HCl and other acids is often written to show the association of the proton with a water molecule to form the hydronium ion [Pg.169]

We begin our study of chemical reactions by examining some of the basic principles of acid-base chemistry. There are several reasons for doing this [Pg.101]

Acid-base reactions also allow us to examine important ideas about the relationship between the structures of molecules and their reactivity and to see how certain thermodynamic parameters can be used to predict how much of the product will be formed when a reaction reaches equilibrium. Acid-base reactions also provide an illustration of the important role solvents play in chemical reactions. They even give us a brief introduction to organic synthesis. Finally, acid-base chemistry is something that you will find familiar because of your studies in general chemistry. We begin, therefore, with a brief review. [Pg.101]

Two classes of acid-base reactions are fundamental in organic chemistry Brpnsted-Lowry and Lewis acid-base reactions. We start our discussion with Brpnsted-Lowry acid-base reactions. [Pg.101]

Hydrogen chloride (HCI), in its pure form, is a gas. When HCI gas is bubbled into water, the following reaction occurs. [Pg.101]

The color of hydrangea flowers depends in part on the relative acidity of their soil. [Pg.101]

Fran s junk-food breakfast and her worrying about the exam combine to give her an annoying case of acid indigestion, which she calms by drinking some baking soda mixed with water. The baking soda contains a base that neutralizes some of her excess stomach acid. [Pg.159]

After taking the exam, Fran feels happy and confident. All those hours working problems, reviewing the learning objectives, and participating in class really paid off. Now she s ready for some lunch. Before eating, she washes her hands with soap made from the reaction of a strong base and animal fat. One of the reasons the soap is slippery is because all bases feel slippery on the skin. [Pg.159]

Fran chooses salad with a piece of lean meat on top for lunch. Like all acids, the vinegar in her salad dressing tastes sour. Ffer stomach produces just enough additional acid to start the digestion of the protein ftom the meat. [Pg.159]

Read on to learn more about the acids and bases that are important in Fran s life and your own what they are, how to construa their names and recognize The vinegar in salad dressing their formulas, and how they react with each other. tastes sour, as do all acids. [Pg.159]

The presentation of information in this chapter assumes that you can already perform the tasks listed below. You can test your readiness to proceed by answering the Review Questions at the end of the chapter. This might also be a good time to read the Chapter Objectives, which precede the Review Questions. [Pg.159]

The overall process can be subdivided into three partial processes (i) the electrochemical reactions, (ii) the acid-base reactions and (iii) the precipitation reactions (Drevet and Benhayoune, 2012). [Pg.147]

The electrochemical reactions occur at the electrode-electrolyte interface and involve the following reactions for calcium nitrate tetrahydrate and ammonium dihydrogen phosphate [Pg.147]

If the pH of the solution is acidic, reduction of the proton can take place according to [Pg.147]

The cathodic reduction reaction (5.2b) increases the pH value of the solution at the cathode-electrolyte interface and gives rise to acid-base reactions. [Pg.147]

Acids and bases are as familiar as aspirin and milk of magnesia although many people do not know their chemical names— acetylsalicylic acid (aspirin) and magnesium hydroxide (milk of magnesia). In addition to being the basis of many medicinal and household products, acid-base chemistry is important in industrial processes and essential in sustaining biological systems. Before we can discuss acid-base reactions, we need to know more about acids and bases themselves. [Pg.129]

In Section 2.7 we defined acids as substances that ionize in water to produce ions and bases as substances that ionize in water to produce OH ions. These definitions were formulated in the late nineteenth century by the Swedish chemist [Pg.129]

Svante Arrhenius to classify substances whose properties in aqueous solutions [Pg.130]

Arrhenius s definitions of acids and bases are limited in that they apply only to aqueous solutions. Broader definitions were proposed by the Danish chemist Johannes Brpnsted in 1932 a Br0nsted acid is a proton donor, and a Br0nsted base is a proton acceptor. Note that Brpnsted s definitions do not require acids and bases to be in aqueous solution. [Pg.130]

Hydrochloric acid is a Brpnsted acid because it donates a proton in water HCl(fl ) ------------------------n iaq) + Cr(fl ) [Pg.130]

The first several chapters of any organic chemistry textbook focus on the structure of molecules how atoms connect to form bonds, how we draw those connections, the problems with our drawing methods, how we name molecules, what molecules look like in 3D, how molecules twist and bend in space, and so on. Only after gaining a clear understanding of stracture do we move on to reactions. But there seems to be one exception acid-base chemistry. [Pg.53]

Acid-base chemistry is typically covered in one of the first few chapters of an organic chemistry textbook, yet it might seem to belong better in the later chapters on reactions. There is an important reason why acid-base chemistry is taught so early on in your course. By understanding this reason, you will have a better perspective of why acid-base chemistry is so incredibly important. [Pg.53]

To appreciate the reason for teaching acid-base chemistry early in the course, we need to first have a very simple understanding of what acid-base chemistry is aU about. Let s summarize with a simple equation [Pg.53]

In the equation above, we see an acid (HA) on the left side of the equilibrium, and the conjugate base (A ) on the right side. HA is an acid by virtue of the fact that it has a proton (H+) to give. A is a base by virtue of the fact that it wants to take its proton back (acids give protons and bases take protons). Since A is the base that we get when we deprotonate HA, we call A the conjugate base of HA. [Pg.53]

So you only need one skill to completely master acid-base chemistry you need to be able to look at a negative charge and determine how stable that negative charge is. If you can do that, then acid-base chemistry will be a breeze for you. If you cannot determine charge stability, then you will have problems even after you finish acid-base chemistry. To predict reactions, you need to know what kind of charges are stable and what kind of charges are not stable. [Pg.53]

The stabilities of species in the system S-Oj-HjO have been studied by numerous researchers. Much of this work has been summarized or critiqued by Garrels and Naeser (1958), Boulegue and Michard (1979), Morse et al. (1987), Schoonen and Barnes (1988), and Williamson and Rimstidt (1992). Thermodynamic data for some substances in the system S-O2-H2O are given in Table A 12.3 in the chapter appendix. The many aqueous sulfur species that are either thermodynamically stable or are important metastably are shown in Fig. 12.14. Acid-base reactions among the sulfur species are generally rapid and reversible. The redox reactions, however, may be fast and reversible (e.g., H2S/SJ see Boulegue and Michard 1979) or more often irreversible in the absence of bacterial activity (e.g., SO4 reduction to H2S). [Pg.445]

Dissociation constants and reaction enthalpy data for the stable and most important metastable sulfur species are summarized in Table 12.4. The log K values in Table 12.4 indicate the pH at which the acid and conjugate base have equal concentrations. Bisulfate (HSO4) is a relatively strong acid. [Pg.445]

Whittemore and D. Langmuir, The solubility of ferric oxyhydroxides in natural waters. Copyright 1975 by Ground Water Publishing Company. Used by permission. [Pg.446]

The value of T (H2S) has been measured by many researchers. Millero (Morse et al. 1987) proposes the temperature function [Pg.446]

The second dissociation constant A 2(H2S) is poorly known. Reported constants range from pA 2 = 12.44 to 17.1 (cf. Morse et al. 1987). The most reliable value is probably that of Schoonen and Barnes (1988), who propose pKi - 18.51 0.56 at 20°C based on extrapolation of thermodynamic data for aqueous polysulfide species. (See also Williamson and Rimstidt 1992.) [Pg.447]

SAMPLE PROBLEM 4.3 Predicting Whether a Precipitation Reaction [Pg.117]

Problem Predict whether a reaction occurs when each of the following pairs of solutions are mixed. If a reaction does occur, write balanced molecular, total ionic, and net ionic equations, and identify the spectator ions. [Pg.117]

Plan For each pair of solutions, we note the ions present in the reactants, write the cation-anion combinations, and refer to Table 4.1 to see if any are insoluble. For the molecular equation, we predict the products. For the total ionic equation, we write the soluble compounds as separate ions. For the net ionic equation, we eliminate the spectator ions. Solution (a) In addition to the reactants, the two other ion combinations are strontium sulfate and sodium nitrate. Table 4.1 shows that strontium sulfate is insoluble, so a reaction does occur. Writing the molecular equation [Pg.117]

FOLLOW-UP PROBLEM 4.3 Predict whether a reaction occurs, and write balanced total and net ionic equations [Pg.117]

Precipitation reactions invoive the formation of an insoiuble ionic compound from two solubie ones. They occur because electrostatic attractions among oertain pairs of solvated ions are strong enough to cause their removai from solution. Suoh reaotions can be predicted by noting whether any possible ion oombinations are insoluble, based on a set of solubility ruies. [Pg.117]

Step 1 The balanced molecular equation for this reaction is [Pg.101]

Step 2 To write the ionic equation, the soluble compounds are shown as dissociated ions [Pg.101]

Step 3 Canceling the spectator ions (K and NO ) on each side of the equation, we obtain the net ionic equation [Pg.101]

Step 4 Note that because we balanced the molecular equation first, the net ionic equation is balanced as to the number of atoms on each side and the number of positive (+6) and negative (-6) charges on the left-hand side is the same. [Pg.101]

Is an insoluble ionic compound, such as AgCI, a strong electrolyte or a weak electrolyte [Pg.161]

Silver chloride, AgCl, is an ionic compound with very low solubility in water. When AgCl dissolves in water, it is 100% dissociated into Ag and CP ions there are no AgCl ion pairs. If we focus only on the degree of dissociation, then AgCl is a strong electrolyte. [Pg.161]

A strong electrolyte may be defined in more practical terms as a substance that, when dissolved in water, gives a solution that is a good conductor of electricity. Because AgCl has very low solubility in water, approximately 1 x 10 moles per liter, a solution of AgCl is not a good conductor of electricity. [Pg.161]

Some chemists would argue that AgCl is a strong electrolyte (because it is 100% dissociated in aqueous solution) but some may argue that it is a weak electrolyte (because an aqueous solution of AgCl is not a good conductor of electricity). [Pg.161]

Ideas about adds and bases (or alkalis) date back to ancient times. The word acid is derived from the Latin acidus (sour). Alkali (base) comes from the Arabic al-qali, referring to the ashes of certain plants from which alkaline substances can be extracted. The acid-base concept is a major theme in the history of chemistry. In this section, we emphasize the view proposed by Svante Arrhenius in 1884 but also introduce a more modern theory proposed in 1923 by Thomas Lowry and by Johannes Bronsted. [Pg.161]

A. (The gas phase estimate is about 100 picoseconds for A at 1 atm pressure.) This suggests tliat tire great majority of fast bimolecular processes, e.g., ionic associations, acid-base reactions, metal complexations and ligand-enzyme binding reactions, as well as many slower reactions that are rate limited by a transition state barrier can be conveniently studied with fast transient metliods. [Pg.2948]

Product acids and bases such as those formed in this process are termed conjugate acids and conjugate bases. Thus, all acid-base reactions can be written as... [Pg.85]

This is an acid-base reaction, in which the base is the oxide ion (p. 89) the acidic oxide SiOj displaces the weaker acidic oxide CO2 in the fused mixture. But in aqueous solution, where the 0 ion cannot function as a strong basefp. 89),carbon dioxide displaces silica, which, therefore, precipitates when the gas is passed through the aqueous silicate solution. In a fused mixture of silica and a nitrate or phosphate, the silica again displaces the weaker acidic oxides N2O5 and P4OJ0 ... [Pg.187]

In Section 1 9 we introduced curved arrows as a tool to systematically generate resonance structures by moving electrons The mam use of curved arrows however is to show the bonding changes that take place in chemical reactions The acid-base reactions to be discussed in Sections 1 12-1 17 furnish numer ous examples of this and deserve some preliminary comment... [Pg.34]

Consider first the case of adding a strong acid such as HBr to water The equation for the Brpnsted acid-base reaction that occurs between them is... [Pg.43]

Analyzing acid-base reactions according to the Brpnsted-Lowry picture provides yet another benefit Table 1 7 which lists acids according to their strength m descending... [Pg.44]

Clearly the two reactions are analogous and demonstrate that the reaction between hydroxide ion and hydrogen bromide is simultaneously a Brpnsted acid-base reaction and a Lewis acid Lewis base reaction Br0nsted acid-base reactions constitute a sub category of Lewis acid Lewis base reactions... [Pg.46]

The position of equilibrium m an acid-base reaction lies to the side of the weaker acid... [Pg.50]

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... [Pg.50]

Although essentially inert m acid-base reactions alkanes do participate m oxidation-reduction reactions as the compound that undergoes oxidation Burning m air (combus tion) IS the best known and most important example Combustion of hydrocarbons is exothermic and gives carbon dioxide and water as the products... [Pg.83]

Our first three chapters established some fundamental principles concerning the structure of organic molecules and introduced the connection between structure and reactivity with a review of acid-base reactions In this chapter we explore structure and reactivity m more detail by developing two concepts functional groups and reaction mechanisms A functional group is the atom or group m a molecule most respon sible for the reaction the compound undergoes under a prescribed set of conditions How the structure of the reactant is transformed to that of the product is what we mean by the reaction mechanism... [Pg.142]

The first step of this new mechanism is exactly the same as that seen earlier for the reaction of tert butyl alcohol with hydrogen chloride—formation of an alkyloxonmm ion by proton transfer from the hydrogen halide to the alcohol Like the earlier exam pie this IS a rapid reversible Brpnsted acid-base reaction... [Pg.164]

The electrophilic character of boron is again evident when we consider the oxida tion of organoboranes In the oxidation phase of the hydroboration-oxidation sequence as presented m Figure 6 11 the conjugate base of hydrogen peroxide attacks boron Hydroperoxide ion is formed m an acid-base reaction m step 1 and attacks boron m step 2 The empty 2p orbital of boron makes it electrophilic and permits nucleophilic reagents such as HOO to add to it... [Pg.254]

The second stage is a Brpnsted acid-base reaction and is fast... [Pg.337]

Step 3 This step is a fast acid base reaction that follows the nucleophilic substitution Water acts as a base to remove a proton from the alkyloxonium ion to give the observed product of the reaction tert butyl alcohol... [Pg.340]

Acetylenic Grignard reagents of the type RC CMgBr are prepared not from an acetylenic halide but by an acid-base reaction in which a Grignard reagent abstracts a proton from a terminal aUcyne... [Pg.598]

Amides are sometimes prepared directly from carboxylic acids and amines by a two step process The first step is an acid-base reaction m which the acid and the amine combine to form an ammonium carboxylate salt On heating the ammonium carboxy late salt loses water to form an amide... [Pg.860]

The acid-base reactions that occur after the amide bond is broken make the overall hydrolysis irreversible m both cases The amine product is protonated m acid the car boxylic acid is deprotonated m base... [Pg.863]

Most of the reactions of ester enolates described so far have centered on stabilized eno lates derived from 1 3 dicarbonyl compounds such as diethyl malonate and ethyl ace toacetate Although the synthetic value of these and related stabilized enolates is clear chemists have long been interested m extending the usefulness of nonstabilized enolates derived from simple esters Consider the deprotonation of an ester as represented by the acid—base reaction... [Pg.902]

Amines like ammonia are weak bases They are however the strongest uncharged bases found m significant quantities under physiological conditions Amines are usually the bases involved m biological acid-base reactions they are often the nucleophiles m biological nucleophilic substitutions... [Pg.913]

In an acid-base reaction, the reaction unit is the proton. For an acid, the number of reaction units is given by the number of protons that can be donated to the base and for a base, the number of reaction units is the number of protons that the base can accept from the acid. In the reaction between H3PO4 and NaOH, for example, the weak acid H3PO4 can donate all three of its protons to NaOH, whereas the strong base NaOH can accept one proton. Thus, we write... [Pg.22]

Several types of reactions are commonly used in analytical procedures, either in preparing samples for analysis or during the analysis itself. The most important of these are precipitation reactions, acid-base reactions, complexation reactions, and oxidation-reduction reactions. In this section we review these reactions and their equilibrium constant expressions. [Pg.139]

The most important types of reactions are precipitation reactions, acid-base reactions, metal-ligand complexation reactions, and redox reactions. In a precipitation reaction two or more soluble species combine to produce an insoluble product called a precipitate. The equilibrium properties of a precipitation reaction are described by a solubility product. [Pg.175]

Acid-base reactions occur when an acid donates a proton to a base. The equilibrium position of an acid-base reaction is described using either the dissociation constant for the acid, fQ, or the dissociation constant for the base, K, . The product of and Kb for an acid and its conjugate base is K (water s dissociation constant). [Pg.175]

A titration in which the reaction between the analyte and titrant is an acid—base reaction. [Pg.278]

Quantitative Calculations In acid-base titrimetry the quantitative relationship between the analyte and the titrant is determined by the stoichiometry of the relevant reactions. As outlined in Section 2C, stoichiometric calculations may be simplified by focusing on appropriate conservation principles. In an acid-base reaction the number of protons transferred between the acid and base is conserved thus... [Pg.304]

Determination of Equilibrium Constants Another important application of molecular absorption is the determination of equilibrium constants. Let s consider, as a simple example, an acid-base reaction of the general form... [Pg.407]

In developing this treatment for determining equilibrium constants, we have considered a relatively simple system in which the absorbance of HIn and Im were easily measured, and for which it is easy to determine the concentration of H3O+. In addition to acid-base reactions, the same approach can be applied to any reaction of the general form... [Pg.409]

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