Jul 5, 2005 — Acid-Base Theory. Chapter Outline. 5.1 Acids and Bases. A comparison of the Arrhenius, Brønsted-Lowry, and. Lewis theories of acids and
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Richard F. Daley and Sally J. Daley www.ochem4free.com Organic Chemistry Chapter 5 Acid-Base Theory 5.1 Acids and Bases 209 5.2 Acid and Base Strength 215 5.3 Hard and Soft Acids and Bases 222 5.4 Organic Acids and Bases 226 5.5 Relative Acidity and Basicity 231 5.6 Substituent Effects on Acidity and Basicity 235 Key Ideas from Chapter 5 238
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Organic Chemistry – Ch 5 205 Daley & Daley Copyright 1996-2005 by Richard F. Daley & Sally J. Daley All Rights Reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, recording, or otherwise, without the prior written permission of the copyright holder. www.ochem4free.com 5 July 2005
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Organic Chemistry – Ch 5 206 Daley & Daley Chapter 5 Acid-Base Theory Chapter Outline 5.1 Acids and Bases A comparison of the Arrhenius, Brønsted-Lowry, and Lewis theories of acids and bases 5.2 Acid and Base Strength A review of pH and Ka 5.3 Hard and Soft Acids and Bases An introduction to hard and soft acid-base theory 5.4 Organic Acids and Bases Molecular characteristics of organic acids and bases 5.5 Relative Acidity and Basicity Estimating relative acidity and basicity 5.6 Substituent Effects on Acidity and Basicity Inductive effects on acid and base strength www.ochem4free.com 5 July 2005
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Organic Chemistry – Ch 5 207 Daley & Daley Objectives Be familiar with the Arrhenius, Brønsted-Lowry, and Lewis theories of acids and bases Recognize the orbitals that are involved in an acid-base reaction Know the relationship between acid strength and the value of pKa Understand the relationship between polarizability and the hardness or softness of an acid or base Predict the stability of a chemical bond using the hard-soft acid base theory Recognize whether an organic functional group is an acid or a base Predict the relative acid or base strength of two organic compounds Understand how the presence of a particular functional group affects the acid or base strength of another functional group I hope no body will offer to dispute whether an Acid has points or no, seeing every ones experience does demonstrate it, they need but to taste an Acid to be satisfied of it, for it pricks the tongue like anything keen, and finely cut – An Alkali is a terrestrous and solid matter, whose pores are figured after such a manner that the Acid points entering them do strike and divide whatsoever opposes their motion. ŠNicholas Lemery “A Course in Chymistry” London (1686) A scof as a you work with chemical reactions in organic hemistry, you will find that you can classify nearly all themcid-base reactions. The key to understanding organic chemical reactions is knowledge of acids and bases. When considering a reaction, you need to ask three questions: Where’s the acid? Where’s the base? How can the acid react with the base? The goal for this chapter is to introduce you to ways that answer these questions. Whether a molecule acts as an acid or a base in a chemical reaction largely depends on its characteristics. There are three www.ochem4free.com 5 July 2005
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Organic Chemistry – Ch 5 208 Daley & Daley significant molecular characteristics that affect acidity and basicity. The most important is the compound’s primary functional group. A second factor is the inductive effect caused by the presence of additional functional groups. A third is the delocalization, or resonance effects, of the electrons in a molecule. Showing Charges on Atoms When you learned to write ions in your introductory chemistry course, you learned to put the charges after the formula of the ion. For example, you wrote the hydroxide ion as OH-. In organic chemistry it is important to know which atom in an ion bears the charge. For example, the oxygen in the hydroxide ion has the negative charge. In this book the hydroxide ion is written as -OH to remind you that the oxygen has the negative charge. Other examples of familiar ions written in this manner are NH4, -CH3, and NO3-. For these three ions, you know immediately that the charges are on N, C, and O respectively. 5.1 Acids and Bases Three major definitions of acids and bases have influenced the thinking of chemists. In 1884, Svante Arrhenius formulated the first of these definitions. Then, in 1923, independently of each other, Johannes N. Brønsted and Thomas M. Lowry developed the second. The third definition grew from Gilbert Newton Lewis’s theory of covalent bonding, which he proposed in 1916. The first definition, proposed by Svante Arrhenius in his doctoral dissertation, was so revolutionary that he was almost denied his Ph.D. However, in 1903, he received the Nobel Prize in chemistry for his theory. His theory states that a stable ionic compound that is soluble in water will break down, or dissociate, into its component ions. This dissociation, or ionization, of a compound in water, leads to Arrhenius’ definition of an acid and a base. An acid is a substance that, when added to water, increases the concentration of hydronium ions, H3O. Because Arrhenius regarded acid-base reactions as occurring only in water, he frequently called the hydronium ion a hydrogen ion, H. An H ion is a proton, or a hydrogen that is electron-deficient. Thus, a base is a substance that, when added to water, increases the concentration of hydroxide ions, -OH. The following statements summarize his definition. An Arrhenius acid is a source of H ion. An Arrhenius base is a source of -OH ion. www.ochem4free.com 5 July 2005
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Organic Chemistry – Ch 5 209 Daley & Daley The Arrhenius acid-base theory provided a good start toward understanding acid-base chemistry, but it proved much too limited in its scope. Brønsted and Lowry developed a more general acid-base definition than that of Arrhenius. Although they considered reactions other than those that take place in aqueous solutions, they still said acids were molecules that donate a hydrogen ionŠsuch as HCl and H2SO4. However, they broadened the definition of bases to include any compound that accepts a proton. The basis of their acid-base definition is that in a reaction a proton transfers between reactants. Thus, acids involving a transfer of H ions are sometimes called proton acids. According to the Brønsted-Lowry definition, an acid is any molecule or ion that donates a proton to another molecule or ion, and a base is any molecule or ion that receives that proton. The following statements briefly summarize the Brønsted-Lowry definition. A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor. An example of the Brønsted-Lowry definition is the reaction between hydrogen chloride and sodium hydroxide: ProtonProtondonoracceptorHCl + NaOHNaCl + H2O In this reaction, HCl is the acid because it is the source of protons, or hydrogen ions; NaOH is a base because the hydroxide ion is the proton acceptor. The following reactions further illustrate the Brønsted-Lowry acid-base definition. H2SO4 + NH3HSO4 + NH4ProtonProtondonoracceptor HCl + CH3CH2NH2Cl + CH3CH2NH3acceptordonorProtonProton www.ochem4free.com 5 July 2005
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Organic Chemistry – Ch 5 211 Daley & Daley proton transfers. Lewis looks at them from the viewpoint of electron pairs. The two viewpoints mesh when you remember that a proton is a positive hydrogen ion that has no electron, and is thus capable of accepting a pair of electrons. Solved Exercise 5.1 The following compounds can act either as a Brønsted-Lowry acid or a Lewis acid. Show the reactive site in each compound and the structure of the conjugate base that results from a reaction with base A-. Determine whether the compound is a Brønsted-Lowry acid or a Lewis acid. a) CH3OH Solution Both the oxygen and the carbon have full valence shells and both have at least one hydrogen as a source of protons. However, oxygen is much more electronegative than carbon, so a negative charge on oxygen is more stable than a negative charge on carbon. Thus, the OŠH bond is the reactive site and a stronger Brønsted-Lowry acid than is the CŠH bond. CH3OH++ACH3OHAAcidConjugate acidConjugate baseBase b) CH3NH2 Solution Nitrogen is much more electronegative than carbon, so a negative charge on nitrogen is more stable than a negative charge on carbon. Thus, the NŠH bond is a stronger Brønsted-Lowry acid than is the CŠH bond. BaseConjugate baseConjugate acidAcidHACH3NHA++CH3NH2 c) CH3BH2 Solution Because boron is electron deficient with only six electrons in its valence shell, it will react before any bonds to hydrogen are broken. Thus, the boron is the reactive site, and it acts as a Lewis acid. BaseAcidCH3BAHHA+CH3BH2 www.ochem4free.com 5 July 2005
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Organic Chemistry – Ch 5 212 Daley & Daley Because a Lewis acid accepts a pair of electrons, chemists call it an electrophile, from the Greek meaning “lover of electrons.” They call the base a nucleophile, or “lover of nuclei,” because it donates the electrons to a nucleus with an empty orbital. In a chemical reaction, a nucleophile seeks a nucleus, or a positive charge, and an electrophile seeks electrons, or a negative charge. Fundamental to organic chemistry is the fact that nearly all the reactions that you will study are reactions of an acid with a base or, more commonly, of an electrophile with a nucleophile. In the formation of a new chemical bond, an electrophile accepts electrons, and a nucleophile donates electrons. Bond formed(base)(acid)ElectrophileNucleophileŁBA+BA NucleophileElectrophile(acid)(base)Bond formedBFFFNHHH+NHHHBFFF Chemists use a curved arrow () to show electron movement. A curved arrow points from the electron-rich reactant, the base or nucleophile, toward the electron-poor reactant, the acid or electrophile. Rewriting the previous two reactions using a curved arrow shows the movement of electrons. In each reaction, a pair of nonbonding electrons from a nucleophile reacts with an electrophile to form a bond. Curved arrows are introduced in Section 1.13, page 000. BA+BA BFFFNHHH+NHHHBFFF Exercise 5.1 Use curved arrows to write the acid-base reaction of a hydrogen ion with a hydroxide ion. www.ochem4free.com 5 July 2005
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Organic Chemistry – Ch 5 213 Daley & Daley Acids and Bases versus Electrophiles and Nucleophiles In organic chemistry, the terms acid and electrophile are formally synonymous, but informally, they have evolved different shades of meaning. The term acid has come to mean a proton donor and the term electrophile has come to mean an electron pair acceptor. Similarly, the term base has come to mean a proton acceptor, and the term nucleophile has come to mean an electron pair donor. However, from time to time, the dividing line between the two sets of terms becomes fuzzy. For example, chemists may call the same group of atoms a base or a nucleophile depending on the chemical environment of that group. Probably the most useful generalization is that the difference between a base and a nucleophile is in how they react. In organic reactions, a base generally reacts with a proton, and a nucleophile generally reacts with a positively charged or electron-deficient carbon. An electron-deficient carbon is a carbon with an unfilled octet in its valence shell. All chemical reactions involve orbital interactions. The orbital description of a reaction can help you understand how chemical reactions occur. As you study the various reactions presented in this book, think about the orbitals involved in the reactions. Figure 5.1 is a molecular orbital picture of ammonia reacting with boron trifluoride to form a new bond. Ammonia is a base with a pair of nonbonding electrons. The nitrogen of ammonia is sp3 hybridized. Boron trifluoride is an acid with an incomplete octet of electrons. The boron is sp2 hybridized with an empty p orbital. The reaction occurs when an sp3 orbital of ammonia overlaps with the empty p orbital of boron trifluoride. In the process, the boron becomes sp3 hybridized. With this overlap the two molecules form a new bond. +BFFFNHHHBFFFNHHH Figure 5.1. The orbitals involved in the acid-base reaction of NH3 and BF3. Exercise 5.2 Show the orbitals involved in the acid-base reaction of a hydrogen ion with a hydroxide ion. Being able to identify an acid or base is important. Of equal importance is the ability to recognize how the structure of that acid or base affects its strength. The rest of this chapter is devoted to helping you acquire the tools to do so. With these tools, you can predict the outcome of chemical reactions. Much of the rest of the material in this book depends on your ability to recognize acids and bases and their relative strengths. www.ochem4free.com 5 July 2005
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Organic Chemistry – Ch 5 214 Daley & Daley 5.2 Acid and Base Strength The strength of a Brønsted-Lowry acid or base depends on the extent to which it ionizes in water. Although there are numerous solvents besides water, chemists discuss acid and base strength in relation to water because they use it so widely as a solvent. Chemists use the autoionization of pure water to determine the values for the concentrations of acidic and basic solutions. Autoionization is the reaction of two molecules of water with each other to give a hydronium ion, H3O, and a hydroxide ion, -OH. Autoionization is a process by which one molecule of a compound reacts with another molecule of the same compound in an acid-base reaction. H3O + OHH2O + H2O For this reaction, the amount of autoionization is extremely slightŠat 25oC, it is 10Œ7 M (moles/liter). The concentrations of H3O and -OH are equal; that is, both measure 10Œ7 M. Chemists call this a neutral solution. If you add a compound that is more acidic than water, you increase the concentration of H3O ions and make the solution acidic. If you add a compound that is more basic than water, you increase the concentration of -OH ions and make the solution basic. The product of the H3O and -OH concentrations in water is equal to 10Œ14 and is a constant, Kw. Chemists define Kw with the following equation. Kw = [H3O][ -OH] = 1.00 x 10Œ14 Because the concentrations of H3O and -OH are equal in a neutral solution, you can easily calculate the concentration of both: When performing a concentration calculation, replace the chemical species listed within the brackets, [-OH], for example, with that species’ molar concentration. [H3O] = [-OH] = 1.00 x 10Œ7 M Because the product of the two concentrations is a constant, Kw, when one concentration increases, the other must decrease. For example, if you add -OH ions to water the concentration of the H3O decreases by whatever amount is necessary for the product of the two concentrations to still equal 10Œ14. Because the hydronium ion concentrations can span a very wide range of values, from greater than 1 M down to less than 10Œ14 M, chemists measure the concentration of H3O on a logarithmic scale called pH. The pH values give the hydronium ion concentration of a solution. Therefore, measuring the pH of a solution is a means of quantifying the acidity of that solution. Chemists define this www.ochem4free.com 5 July 2005
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