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- Herausgeber: John Wiley & Sons
- Kategorie: Wissenschaft und neue Technologien
- Sprache: Englisch
KEYNOTES IN Organic Chemistry
KEYNOTES IN Organic Chemistry
SECOND EDITION
This concise and accessible textbook provides notes for students studying chemistry and related courses at undergraduate level, covering core organic chemistry in a format ideal for learning and rapid revision. The material, with an emphasis on pictorial presentation, is organised to provide an overview of the essentials of functional group chemistry and reactivity, leading the student to a solid understanding of the basics of organic chemistry.
This revised and updated second edition of Keynotes in Organic Chemistry includes:
- new margin notes to emphasise links between different topics,
- colour diagrams to clarify aspects of reaction mechanisms and illustrate key points, and
- a new keyword glossary.
In addition, the structured presentation provides an invaluable framework to facilitate the rapid learning, understanding and recall of critical concepts, facts and definitions. Worked examples and questions are included at the end of each chapter to test the reader’s understanding.
Reviews of the First Edition
“ …this text provides an outline of what should be known and understood, including fundamental concepts and mechanisms.”
Journal of Chemical Education, 2004
“ Despite the book’s small size, each chapter is thorough, with coverage of all important reactions found at first-year level... ideal for the first-year student wishing to revise… and priced and designed appropriately.”
The Times Higher Education Supplement, 2004
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Veröffentlichungsjahr: 2013
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Contents
Cover
Title Page
Copyright
Preface
Chapter 1: Structure and Bonding
1.1 Ionic versus Covalent Bonds
1.2 The Octet Rule
1.3 Formal Charge
1.4 Sigma (σ–) and pi (π–) Bonds
1.5 Hybridisation
1.6 Inductive Effects, Hyperconjugation and Mesomeric Effects
1.7 Acidity and Basicity
Worked Example
Problems
Chapter 2: Functional Groups, Nomenclature and Drawing Organic Compounds
2.1 Functional Groups
2.2 Alkyl and Aryl Groups
2.3 Alkyl Substitution
2.4 Naming Carbon Chains
2.5 Drawing Organic Structures
Worked Example
Problems
Chapter 3: Stereochemistry
3.1 Isomerism
3.2 Conformational Isomers
3.3 Configurational Isomers
Worked Example
Problems
Chapter 4: Reactivity and Mechanism
4.1 Reactive Intermediates: ions versus Radicals
4.2 Nucleophiles and Electrophiles
4.3 Carbocations, Carbanions and Carbon Radicals
4.4 Steric Effects
4.5 Oxidation Levels
4.6 General types of Reaction
4.7 Ions versus Radicals
4.8 Reaction Selectivity
4.9 Reaction Thermodynamics and Kinetics
4.10 Orbital Overlap and Energy
4.11 Guidelines for Drawing Reaction Mechanisms
Worked Example
Problems
Chapter 5: Halogenoalkanes
5.1 Structure
5.2 Preparation
5.3 Reactions
Worked Example
Problems
Chapter 6: Alkenes and Alkynes
6.1 Structure
6.2 Alkenes
6.3 Alkynes
Worked Example
Problems
Chapter 7: Benzenes
7.1 Structure
7.2 Reactions
7.3 Reactivity of Substituted Benzenes
7.4 Nucleophilic Aromatic Substitution (the SNAr Mechanism)
7.5 The Formation of Benzyne
7.6 Transformation of Side Chains
7.7 Reduction of the Benzene Ring
7.8 The Synthesis of Substituted Benzenes
7.9 Electrophilic Substitution of Naphthalene
7.10 Electrophilic Substitution of Pyridine
7.11 Electrophilic Substitution of Pyrrole, Furan and Thiophene
Worked Example
Problems
Chapter 8: Carbonyl Compounds: Aldehydes and Ketones
8.1 Structure
8.2 Reactivity
8.3 Nucleophilic Addition Reactions
8.4 α-Substitution Reactions
8.5 Carbonyl-Carbonyl Condensation Reactions
Worked Example
Problems
Chapter 9: Carbonyl Compounds: Carboxylic Acids and Derivatives
9.1 Structure
9.2 Reactivity
9.3 Nucleophilic Acyl Substitution Reactions
9.4 Nucleophilic Substitution Reactions of Carboxylic Acids
9.5 Nucleophilic Substitution Reactions of Acid Chlorides
9.6 Nucleophilic Substitution Reactions of Acid Anhydrides
9.7 Nucleophilic Substitution Reactions of Esters
9.8 Nucleophilic Substitution and Reduction Reactions of Amides
9.9 Nucleophilic Addition Reactions of Nitriles
9.10 α-Substitution Reactions of Carboxylic Acids
9.11 Carbonyl-Carbonyl Condensation Reactions
9.12 A Summary of Carbonyl Reactivity
Worked Example
Problems
Chapter 10: Spectroscopy
10.1 Mass Spectrometry (MS)
10.2 The Electromagnetic Spectrum
10.3 Ultraviolet (UV) Spectroscopy
10.4 Infrared (IR) Spectroscopy
10.5 Nuclear Magnetic Resonance (NMR) Spectroscopy
Worked Example
Problems
Chapter 11: Natural Products and Synthetic Polymers
11.1 Carbohydrates
11.2 Lipids
11.3 Amino Acids, Peptides and Proteins
11.4 Nucleic Acids
11.5 Synthetic Polymers
Worked Example
Problems
Appendix 1: Bond Dissociation Enthalpies
Appendix 2: Bond Lengths
Appendix 3: Approximate pKa Values (Relative to Water)
Appendix 4: Useful Abbreviations
Appendix 5: Infrared Absorptions
Appendix 6: Approximate NMR Chemical Shifts
Appendix 7: Reaction Summaries
Appendix 8: Glossary
Further Reading
Outline Answers
Index
This edition first published 2014
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Library of Congress Cataloging-in-Publication Data
Parsons, A. F.
Keynotes in organic chemistry / Andrew Parsons. – Second edition.
pages cm.
Includes bibliographical references and index.
ISBN 978-1-119-99915-7(hardback) – ISBN 978-1-119-99914-0 (paperback) 1. Chemistry, Organic–Outlines, syllabi, etc. I. Title.
QD256.5.P35 2014
547–dc23
2013024694
A catalogue record for this book is available from the British Library.
HB ISBN: 9781119999157
PB ISBN: 9781119999140
Preface
With the advent of modularisation and an ever-increasing number of examinations, there is a growing need for concise revision notes that encapsulate the key points of a subject in a meaningful fashion. This keynote revision guide provides concise organic chemistry notes for first year students studying chemistry and related courses (including biochemistry) in the UK. The text will also be appropriate for students on similar courses in other countries.
An emphasis is placed on presenting the material pictorially (pictures speak louder than words); hence, there are relatively few paragraphs of text but numerous diagrams. These are annotated with key phrases that summarise important concepts/key information and bullet points are included to concisely highlight key principles and definitions.
The material is organised to provide a structured programme of revision. Fundamental concepts, such as structure and bonding, functional group identification and stereochemistry are introduced in the first three chapters. An important chapter on reactivity and mechanism is included to provide a short overview of the basic principles of organic reactions. The aim here is to provide the reader with a summary of the ‘key tools’ which are necessary for understanding the following chapters and an important emphasis is placed on organisation of material based on reaction mechanism. Thus, an overview of general reaction pathways/mechanisms (such as substitution and addition) is included and these mechanisms are revisited in more detail in the following chapters. Chapters 5–10 are treated essentially as ‘case studies’, reviewing the chemistry of the most important functional groups. Halogenoalkanes are discussed first and as these compounds undergo elimination reactions this is followed by the (electrophilic addition) reactions of alkenes and alkynes. This leads on to the contrasting (electrophilic substitution) reactivity of benzene and derivatives in Chapter 7, while the rich chemistry of carbonyl compounds is divided into Chapters 8 and 9. This division is made on the basis of the different reactivity (addition versus substitution) of aldehydes/ketones and carboxylic acid derivatives to nucleophiles. A chapter is included to revise the importance of spectroscopy in structure elucidation and, finally, the structure and reactivity of a number of important natural products and synthetic polymers is highlighted in Chapter 11. Worked examples and questions are included at the end of each chapter to test the reader's understanding, and outline answers are provided for all of the questions. Tables of useful physical data, reaction summaries and a glossary are included in appendices at the back of the book.
New to this Edition
A number of additions have been made to this edition to reflect the feedback from students and lecturers:
A second colour is used to clarify some of the diagrams, particularly the mechanistic aspectsReference notes are added in the margin to help the reader find information and to emphasise links between different topicsDiagrams are included in the introductory key pointsections for each chapterAdditional end-of-chapter problems (with outline answers) are includedA worked example is included at the end of each chapterThe information in the appendices has been expanded, including reaction summaries and a glossaryAcknowledgements
There are numerous people I would like to thank for their help with this project. This includes many students and colleagues at York. Their constructive comments were invaluable. I would also like to thank my family for their support and patience throughout this project. Finally, I would like to thank Paul Deards and Sarah Tilley from Wiley, for all their help in progressing the second edition.
Dr Andrew F. Parsons
2013
1
Structure and Bonding
Key point. Organic chemistry is the study of carbon compounds. Ionic bonds involve elements gaining or losing electrons but the carbon atom is able to form four covalent bonds by sharing the four electrons in its outer shell. Single (C–C), double (C=C) or triple bonds (CC) to carbon are possible. When carbon is bonded to a different element, the electrons are not shared equally, as electronegative atoms (or groups) attract the electron density whereas electropositive atoms (or groups) repel the electron density. An understanding of the electron-withdrawing or -donating ability of atoms, or a group of atoms, can be used to predict whether an organic compound is a good acid or base.
1.1 Ionic versus Covalent Bonds
Ionic bonds are formed between molecules with opposite charges. The negatively charged anion will electrostatically attract the positively charged cation. This is present in (inorganic) salts.Covalent bonds are formed when a pair of electrons is shared between two atoms. A single line represents the two-electron bond.Coordinate (or dative) bonds are formed when a pair of electrons is shared between two atoms. One atom donates both electrons and a single line or an arrow represents the two-electron bond.The cyclic ether is tetrahydrofuran (THF) and BH3is called borane ( Section 6.2.2.5)
Hydrogen bonds are formed when the partially positive (δ+) hydrogen of one molecule interacts with the partially negative (δ−) heteroatom (e.g. oxygen or nitrogen) of another molecule.Intramolecular hydrogen bonding in carbonyl compounds is discussed in Section 8.4.1
1.2 The Octet Rule
To form organic compounds, the carbon atom shares electrons to give a stable ‘full shell’ electron configuration of eight valence electrons.
Methane is the smallest alkane – alkanes are a family of compounds that contain only C and H atoms linked by single bonds (Section 2.4)
Drawing organic compounds using full structural formulae and other conventions is discussed in Section 2.5
A single bond contains two electrons, a double bond contains four electrons and a triple bond contains six electrons. A lone (or non-bonding) pair of electrons is represented by two dots (• •).
1.3 Formal Charge
Formal positive or negative charges are assigned to atoms, which have an apparent ‘abnormal’ number of bonds.
Nitric acid is used in synthesis to nitrate aromatic compounds such as benzene (Section 7.2.2)
Carbon forms four covalent bonds. When only three covalent bonds are present, the carbon atom can have either a formal negative charge or a formal positive charge.
The stability of carbocations and carbanions is discussed in Section 4.3
Carbanions –three covalent bonds to carbon and a formal negative charge.Carbanions are formed on deprotonation of organic compounds. Deprotonation of a carbonyl compound, at the α-position, forms a carbanion called an enolate ion (Section 8.4.3)
Carbocations –three covalent bonds to carbon and a formal positive charge.Carbocations are intermediates in a number of reactions, including S N1 reactions (Section 5.3.1.2)
1.4 Sigma (σ–) and pi (π–) Bonds
The electrons shared in a covalent bond result from overlap of atomic orbitals to give a new molecular orbital. Electrons in 1s and 2s orbitals combine to give sigma (σ–) bonds.
Molecular orbitals and chemical reactions are discussed in Section 4.10
When two 1s orbitals combine in-phase, this produces a bonding molecular orbital.
When two 1s orbitals combine out-of-phase, this produces an antibonding molecular orbital.
Electrons in p orbitals can combine to give sigma (σ) or pi (π) bonds.
Sigma (σ–) bonds are strong bonds formed by head-on overlap of two atomic orbitals.Pi (π–) bonds are weaker bonds formed by side-on overlap of two p-orbitals.Alkenes have a C C bond containing one strong σ-bond and one weaker π-bond (Section 6.1)
All carbonyl compounds have a C O bond, which contains one strong σ-bond and one weaker π-bond (Section 8.1)
Only σ- or π-bonds are present in organic compounds. All single bonds are σ-bonds while all multiple (double or triple) bonds are composed of one σ-bond and one or two π-bonds.
1.5 Hybridisation
Hund's rule states that when filling up a set of orbitals of the same energy, electrons are added with parallel spins to different orbitals rather than pairing two electrons in one orbital
The ground-state electronic configuration of carbon is 1s22s22px12py1.The six electrons fill up lower energy orbitals before entering higher energy orbitals (Aufbau principle).Each orbital is allowed a maximum of two electrons (Pauli exclusion principle).The two 2p electrons occupy separate orbitals before pairing up (Hund's rule).The carbon atom can mix the 2s and 2p atomic orbitals to form four new hybrid orbitals in a process known as hybridisation.
sp3Hybridisation. For four single σ-bonds – carbon is sp3 hybridised (e.g. in methane, CH4). The orbitals move as far apart as possible, and the lobes point to the corners of a tetrahedron (109.5° bond angle).Alkenes have a C C bond containing one strong σ-bond and one weaker π-bond (Section 6.1)
sp2Hybridisation. For three single σ-bonds and one π-bond – the π-bond requires one p-orbital, and hence the carbon is sp2 hybridised (e.g. in ethene, H2C=CH2). The three sp2-orbitals point to the corners of a triangle (120° bond angle), and the remaining p-orbital is perpendicular to the sp2 plane.All carbonyl compounds have a C O bond, which contains one strong σ-bond and one weaker π-bond (Section 8.1)
Alkynes have a CC bond containing one strong σ-bond and two weaker π-bonds (Section 6.1)
sp Hybridisation. For two single σ-bonds and two π-bonds – the two π-bonds require two p-orbitals, and hence the carbon is sp hybridised (e.g. in ethyne, HCCH). The two sp-orbitals point in the opposite directions (180° bond angle), and the two p-orbitals are perpendicular to the sp plane.For a single C–C or C–O bond, the atoms are sp3 hybridised and the carbon atom(s) is tetrahedral.For a double C=C or C=O bond, the atoms are sp2 hybridised and the carbon atom(s) is trigonal planar.For a triple CC or CN bond, the atoms are sp hybridised and the carbon atom(s) is linear.This compound contains four functional groups, including a phenol. Functional groups are introduced in Section 2.1
The shape of organic molecules is therefore determined by the hybridisation of the atoms.
Functional groups (Section 2.1) that contain π-bonds are generally more reactive as a π-bond is weaker than a σ-bond. The π-bond in an alkene or alkyne is around +250 kJ mol−1, while the σ-bond is around +350 kJ mol−1.
The shorter the bond length, the stronger the bond. For C–H bonds, the greater the ‘s’ character of the carbon orbitals, the shorter the bond length. This is because the electrons are held closer to the nucleus.
A hydrogen atom attached to a CC bond is more acidic than a hydrogen atom attached to a C C bond or a C– C bond; this is explained by the change in hybridisation of the carbon atom that is bonded to the hydrogen atom (Section 1.7.4)
A single C–C σ-bond can undergo free rotation at room temperature, but a π-bond prevents free rotation around a C=C bond. For maximum orbital overlap in a π-bond, the two p-orbitals need to be parallel to one another. Any rotation around the C=C bond will break the π-bond.
Rotation about C–C bonds is discussed in Section 3.2
1.6 Inductive Effects, Hyperconjugation and Mesomeric Effects
1.6.1 Inductive Effects
In a covalent bond between two different atoms, the electrons in the σ-bond are not shared equally. The electrons are attracted towards the most electronegative atom. An arrow drawn above the line representing the covalent bond can show this. (Sometimes an arrow is drawn on the line.) Electrons are pulled in the direction of the arrow.
An inductive effect is the polarisation of electrons through σ-bonds
An alkyl group (R) is formed by removing a hydrogen atom from an alkane (Section 2.2).
An aryl group (Ar) is benzene (typically called phenyl, Ph) or a substituted benzene group (Section 2.2)
The overall polarity of a molecule is determined by the individual bond polarities, formal charges and lone pair contributions and this can be measured by the dipole moment (μ). The larger the dipole moment (often measured in debyes, D), the more polar the compound.
1.6.2 Hyperconjugation
A σ-bond can stabilise a neighbouring carbocation (or positively charged carbon, e.g. R3C+) by donating electrons to the vacant p-orbital. The positive charge is delocalised or ‘spread out’ and this stabilising effect is called resonance.
Hyperconjugation is the donation of electrons from nearby C– H or C– C σ-bonds
The stability of carbocations is discussed in Section 4.3.1
1.6.3 Mesomeric Effects
Whilst inductive effects pull electrons through the σ-bond framework, electrons can also move through the π-bond network. A π-bond can stabilise a negative charge, a positive charge, a lone pair of electrons or an adjacent bond by resonance (i.e. delocalisation or ‘spreading out’ of the electrons). Curly arrows (Section 4.1) are used to represent the movement of π- or non-bonding electrons to give different resonance forms. It is only the electrons, not the nuclei, that move in the resonance forms, and a double-headed arrow is used to show their relationship.
Resonance forms (sometimes called canonical forms) show all possible distributions of electrons in a molecule or an ion
1.6.3.1 Positive Mesomeric Effect
When a π-system donates electrons, the π-system has a positive mesomeric effect, +M.This carbocation is called an allylic cation (see Section 5.3.1.2)
When a lone pair of electrons is donated, the group donating the electrons has a positive mesomeric effect, +M.The OR group is called an alkoxy group (see Section 2.4)
1.6.3.2 Negative Mesomeric Effect
When a π-system accepts electrons, the π-system has a negative mesomeric effect, −M.This anion, formed by deprotonating an aldehyde at the α-position, is called an enolate ion (Section 8.4.3)
The actual structures of the cations or anions lie somewhere between the two resonance forms. All resonance forms must have the same overall charge and obey the same rules of valency.
Functional groups are discussed in Section 2.1
In neutral compounds, there will always be a +M and −M group(s): one group donates (+M) the electrons, the other group(s) accepts the electrons (−M).
An amide, such as RCONH2, also contains both a +M group (NH2) and a −M group (C=O). See Sections 1.7.2 and 9.3.1
All resonance forms are not of the same energy. Generally, the most stable resonance forms have the greatest number of covalent bonds, atoms with a complete valence shell of electrons, and/or an aromatic ring. In phenol (PhOH), for example, the resonance form with the intact aromatic benzene ring is expected to predominate.
Benzene and other aromatic compounds, including phenol, are discussed in Chapter 7
As a rule of thumb, the more resonance structures an anion, cation or neutral π-system can have, the more stable it is.
1.6.3.3 Inductive versus Mesomeric Effects
Mesomeric effects are generally stronger than inductive effects. A +M group is likely to stabilise a cation more effectively than a +I group.
Mesomeric effects can be effective over much longer distances than inductive effects provided that conjugation is present (i.e. alternating single and double bonds). Whereas inductive effects are determined by distance, mesomeric effects are determined by the relative positions of +M and −M groups in a molecule (Section 1.7).
Conjugated enones, containing a C C– C O group, are discussed in Section 8.5.1
1.7 Acidity and Basicity
1.7.1 Acids
An acid is a substance that donates a proton (Brønsted-Lowry). Acidic compounds have low pKa values and are good proton donors as the anions (or conjugate bases), formed on deprotonation, are relatively stable.
Equilibria and equilibrium constants are discussed in Section 4.9.1.1
The pKa value equals the pH of the acid when it is half ionised. At pH's above the pKa the acid (HA) exists predominantly as the conjugate base (A−) in water. At pH's below the pKa it exists predominantly as HA.
The pKa values are influenced by the solvent. Polar solvents will stabilise cations and/or anions by solvation in which the charge is delocalised over the solvent (e.g. by hydrogen-bonding in water).
The influence of solvent polarity on substitution and elimination reactions is discussed in Sections 5.3.1.3 and 5.3.2.3
The more electronegative the atom bearing the negative charge, the more stable the conjugate base (which is negatively charged).
Therefore, F− is more stable than H3C−.
Inductive effects are introduced in Section 1.6.1
The conjugate base can also be stabilised by −I and −M groups which can delocalise the negative charge. (The more ‘spread out’ the negative charge, the more stable it is).
Mesomeric effects are introduced in Section 1.6.3
1.7.1.1 Inductive Effects and Carboxylic Acids
The carboxylate ion (RCO2−) is formed on deprotonation of a carboxylic acid (RCO2H). The anion is stabilised by resonance (i.e. the charge is spread over both oxygen atoms) but can also be stabilised by the R group if this has a −I effect.
The reactions of carboxylic acids are discussed in Chapter 9
The greater the −I effect, the more stable the carboxylate ion (e.g. FCH2CO2− is more stable than BrCH2CO2−) and the more acidic the carboxylic acid (e.g. FCH2CO2H is more acidic than BrCH2CO2H).
1.7.1.2 Inductive and Mesomeric Effects and Phenols
Mesomeric effects can also stabilise positive and negative charges.
On deprotonation of phenol (PhOH) the phenoxide ion (PhO−) is formed. This anion is stabilised by the delocalisation of the negative charge on to the 2-, 4- and 6-positions of the benzene ring.
If −M groups are introduced at the 2-, 4- and/or 6-positions, the anion can be further stabilised by delocalisation through the π-system as the negative charge can be spread onto the −M group. We can use double-headed curly arrows to show this process.Double-headed curly arrows are introduced in Section 4.1
If −M groups are introduced at the 3- and/or 5-positions, the anion cannot be stabilised by delocalisation, as the negative charge cannot be spread onto the −M group. There is no way of using curly arrows to delocalise the charge on to the −M group.If −I groups are introduced on the benzene ring, the effect will depend on their distance from the negative charge. The closer the −I group is to the negative charge, the greater the stabilising effect will be. The order of −I stabilisation is therefore 2-position > 3-position > 4-position.The −M effects are much stronger than −I effects (Section 1.6.3).Examples
Naming substituted benzenes is discussed in Section 2.4
1.7.2 Bases
A base is a substance that accepts a proton (Brønsted-Lowry). Basic compounds are good proton acceptors as the conjugate acids, formed on protonation, are relatively stable. Consequently, strong bases (B: or B−) give conjugate acids (BH+ or BH) with high pKa values.
Equilibria and equilibrium constants are discussed in Section 4.9.1.1
The strength of bases is usually described by the Ka and pKa values of the conjugate acid.
For the use of bases in elimination reactions of halogenoalkanes, see Section 5.3.2
For reactions of bases with carbonyl compounds see Sections 8.4.3 and 9.11
If B is a strong base then BH+ will be relatively stable and not easily deprotonated. BH+ will therefore have a high pKa value.If B is a weak base then BH+ will be relatively unstable and easily deprotonated. BH+ will therefore have a low pKa value.Inductive effects are introduced in Section 1.6.1
The cation can be stabilised by +I and +M groups, which can delocalise the positive charge. (The more ‘spread out’ the positive charge, the more stable it is.)
Mesomeric effects are introduced in Section 1.6.3
1.7.2.1 Inductive Effects and Aliphatic (or Alkyl) Amines
On protonation of amines (e.g. RNH2), ammonium salts are formed.
Aliphatic amines have nitrogen bonded to one or more alkyl groups; aromatic amines have nitrogen bonded to one or more aryl groups
Primary (RNH2), secondary (R2NH) and tertiary (R3N) amines are introduced in Section 2.1
The greater the +I effect of the R group, the greater the electron density at nitrogen and the more basic the amine. The greater the +I effect, the more stable the ammonium ion and the more basic the amine.
Triethylamine (Et3N) is commonly used as a base in organic synthesis (Section 5.2.2)
The pKa values should increase steadily as more +I alkyl groups are introduced on nitrogen. However, the pKa values are determined in water, and the more hydrogen atoms on the positively charged nitrogen, the greater the extent of hydrogen bonding between water and the cation. This solvation leads to the stabilisation of the cations containing N–H bonds.
Hydrogen bonds are introduced in Section 1.1
In organic solvents (which cannot solvate the cation) the order of pKa's is expected to be as follows.
The presence of −I and/or −M groups on nitrogen reduces the basicity and so, for example, primary amides (RCONH2) are poor bases.
Secondary amides (RCONHR) and tertiary amides (RCONR2) are also very weak bases because the nitrogen lone pairs are stabilised by resonance
If ethanamide was protonated on nitrogen, the positive charge could not be stabilised by delocalisation. Protonation therefore occurs on oxygen as the charge can be delocalised on to the nitrogen atom.
Reactions of amides are discussed in Section 9.8
1.7.2.2 Mesomeric Effects and Aryl (or Aromatic) Amines
The lone pair of electrons on the nitrogen atom of aminobenzene (or aniline, PhNH2) can be stabilised by delocalisation of the electrons onto the 2-, 4- and 6-positions of the benzene ring. Aromatic amines are therefore less basic than aliphatic amines.
For the preparation and reactions of aniline (PhNH2), see Section 7.8
If −M groups are introduced at the 2-, 4- and/or 6-positions (but not the 3- or 5-positions) the anion can be further stabilised by delocalisation, as the negative charge can be spread on to the −M group. This reduces the basicity of the amine.If −I groups are introduced on the benzene ring, the order of −I stabilisation is 2-position > 3-position > 4-position. This reduces the basicity of the amine.If +M groups (e.g. OMe) are introduced at the 2-, 4- or 6-position of aminobenzene (PhNH2), then the basicity is increased. This is because the +M group donates electron density to the carbon atom bearing the amine group. Note that the nitrogen atom, not the oxygen atom, is protonated – this is because nitrogen is less electronegative than oxygen and is a better electron donor.For the Pauling electronegativity scale see Section 1.6.1
The OMe group is called a methoxy group (see Section 2.4 for naming organic compounds)
Curly arrows can be used to show the delocalisation of electrons on to the carbon atom bearing the nitrogen.
Curly arrows are introduced in Section 4.1
1.7.3 Lewis Acids and Bases
A heteroatom is any atom that is not carbon or hydrogen
A Lewis acid is any substance that accepts an electron pair in forming a coordinate bond (Section 1.1). Examples include H+, BF3, AlCl3, TiCl4, ZnCl2 and SnCl4. They have unfilled valence shells and so can accept electron pairs.A Lewis base is any substance that donates an electron pair in forming a coordinate bond. Examples include H2O, ROH, RCHO, R2C=O, R3N and R2S. They all have a lone pair(s) of electrons on the heteroatom (O, N or S).Reactions of ketones are discussed in Chapter 8
1.7.4 Basicity and Hybridisation
The greater the ‘s’ character of an orbital, the lower in energy the electrons and the more tightly the electrons are held to the nucleus. The electrons in an sp-orbital are therefore less available for protonation than those in an sp2- or sp3-orbital, and hence the compounds are less basic.
1.7.5 Acidity and Aromaticity
Aromatic compounds are planar, conjugated systems which have 4n + 2 electrons (Hückel's rule) (Section 7.1). If, on deprotonation, the anion is part of an aromatic π-system then the negative charge will be stabilised. Aromaticity will therefore increase the acidity of the compound.
Toluene is a common solvent. Oxidation of the CH3 group is discussed in Section 7.6
Resonance stabilisation of carbanions is introduced in Section 1.6.3
If a lone pair of electrons on a heteroatom is part of an aromatic π-system, then these electrons will not be available for protonation. Aromaticity will therefore decrease the basicity of the compound.
Reactions of aromatic heterocycles, including pyrrole and pyridine are discussed in Sections 7.10 and 7.11
1.7.6 Acid-Base Reactions
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 with a higher pKa value. This means that the product acid and base will be more stable than the starting acid and base.
For a table of pKa values see Appendix 3
Deprotonation of terminal alkynes is discussed in Section 6.3.2.5
For deprotonation of carbonyl compounds to form enolate ions, see Section 8.4.3
Worked Example
Hint: Determine whether the groups attached to the negatively charged carbons in 1 −4 can stabilise the lone pair by I and/or M effects
Hint: Consider a δ+ hydrogen atom bonded to an electronegative atom that, on deprotonation, gives the more stable conjugate base
Hint: Show all the lone pairs in 6 and consider their relative availability. Compare the stability of possible conjugate acids
Answer
Inductive and mesomeric effects (resonance) are discussed in Sections 1.6.1 and 1.6.3
For the preparation and reactions of enolate ions, see Section 8.4.3
Formation of carboxylate ions is discussed in Section 1.7.1
Functional groups are discussed in Section 2.1
Formation of ammonium ions is discussed in Section 1.7.2
Problems
Sections 1.6.1 and 1.6.3
1. Using the I and M notations, identify the electronic effects of the following substituents.
Section 1.6.3
2.
Sections 1.6.3 and 1.7.1
3. Provide explanations for the following statements.
Section 1.7.5
4. Why is cyclopentadiene (pKa 15.5) a stronger acid than cycloheptatriene (pKa ∼ 36)?
Section 1.7.1
5. Which hydrogen atom would you expect to be the most acidic in each of the following compounds?
Section 1.7.2
6. Arrange the following sets of compounds in order of decreasing basicity. Briefly explain your reasoning.
7. For each of the following compounds D −F, identify the most acidic hydrogen atom(s). Briefly explain your reasoning.
Section 1.7.1
8. For each of the following compounds G −I, identify the most basic group. Briefly explain your reasoning.
Section 1.7.2
9. Given the approximate pKa values shown below, for the following acid-base reactions (a)−(e), determine whether the position of the equilibrium lies over to the reactant side or the product side.
Section 1.7.6
AcidpKa valuePhOH9.9H2O15.7CH3COCH320H235NH338H2C=CH244
