PREFACE a comprehensiveset of basic chemistry,wh h will be suitable for und studen chemistry-re course: or cou chemistry as an ancillary subject.The book concentrates on core topics which are most likely to be common to those organic chemistry courses which follow on from a foun- dation or introductory general chemistry course Organic chemistry is a subject which can lead some students to the heights of ecstasy,yet drive others up the wall.Some students 'switch on'to it imme diately,while others can make neither head nor tail of it,no matter how hard they try.Certainly,one of the major problems in studying the subject is the vast amount of material which often has to be covered.Many students blanche at the prospect of having to learn a seemingly endless number of reactions and when it comes to drawing mechanisms and curly arrows,they see only a confusing maze of squiggly lines going everywhere yet nowhere.The concepts of organic reaction mechanisms are often the most difficult to master. These difficulties are often compounded by the fact that current textbooks in organic chemistry are typically over 1200 pages long and can be quite expen- sive to buy. This book attempts to condense the essentials of organic chemistry into a manageable text of 310 pages which is student friendly and which does not cost an arm and a leg.It does this by concentrating purely on the basics of the subject without going into exhaustive detail or repetitive examples. Furthermore,key notes at the start of each topic summarize the essential facts covered and help focus the mind on the essentials. Organic chemistry is a peculiar subject in that it becomes easier as you go along!This might seem an outrageous statement to make,esp pecially to a first- year studenthmotermsth the ru of nmenature. trying to memorize a couple of dozen reactions and making sense of mecha- nisms at the same time.However,these topics are the basics of the subject and once they have been grasped,the overall p icture hecomes clear Understanding the me hanism of how a reaction takes place is particularly crucial in this.It brings a logic to the reactions of the different functional gr This in turn transfor ms a list of ar which makes remembering the aetyuneatedacsiaioaStebl5ep reactions a 'piece of cake'(well,nearly). Once this ha state of affairs has been reached.the relevance of organic etics and biochemistry r suddenly lea off the p organic chemistry leads to a better unders tandi of life chemistry and how the ody works at the molecular level.It also helps in the nding of the molecular mecha volved in dise n to ar ug an be designed o those dis edicinal che A that's not all An unde ic che the ndus- trial che or che ith P chemical nd the cientist trying proces and syn the d fu eco frie
This textbook aims to provide a comprehensive set of basic notes in organic chemistry, which will be suitable for undergraduate students taking chemistry, chemistry-related courses, or courses which involve organic chemistry as an ancillary subject. The book concentrates on core topics which are most likely to be common to those organic chemistry courses which follow on from a foundation or introductory general chemistry course. Organic chemistry is a subject which can lead some students to the heights of ecstasy, yet drive others up the wall. Some students ‘switch on’ to it immediately, while others can make neither head nor tail of it, no matter how hard they try. Certainly, one of the major problems in studying the subject is the vast amount of material which often has to be covered. Many students blanche at the prospect of having to learn a seemingly endless number of reactions, and when it comes to drawing mechanisms and curly arrows, they see only a confusing maze of squiggly lines going everywhere yet nowhere. The concepts of organic reaction mechanisms are often the most difficult to master. These difficulties are often compounded by the fact that current textbooks in organic chemistry are typically over 1200 pages long and can be quite expensive to buy. This book attempts to condense the essentials of organic chemistry into a manageable text of 310 pages which is student friendly and which does not cost an arm and a leg. It does this by concentrating purely on the basics of the subject without going into exhaustive detail or repetitive examples. Furthermore, key notes at the start of each topic summarize the essential facts covered and help focus the mind on the essentials. Organic chemistry is a peculiar subject in that it becomes easier as you go along! This might seem an outrageous statement to make, especially to a firstyear student who is struggling to come to terms with the rules of nomenclature, trying to memorize a couple of dozen reactions and making sense of mechanisms at the same time. However, these topics are the basics of the subject and once they have been grasped, the overall picture becomes clear. Understanding the mechanism of how a reaction takes place is particularly crucial in this. It brings a logic to the reactions of the different functional groups. This in turn transforms a list of apparently unrelated facts into a sensible theme which makes remembering the reactions a ‘piece of cake’ (well, nearly). Once this happy state of affairs has been reached, the relevance of organic chemistry to other subjects such as genetics and biochemistry suddenly leaps off the page. Understanding organic chemistry leads to a better understanding of life chemistry and how the body works at the molecular level. It also helps in the understanding of the molecular mechanisms involved in disease and bodily malfunction, leading in turn to an understanding of how drugs can be designed to cure these disease states – the science of medicinal chemistry. And that’s not all. An understanding of organic chemistry will help the industrial chemist or chemical engineer faced with unexpected side-reactions in a chemical process, and the agro-scientist trying to understand the molecular processes taking place within plants and crops; and it will assist in the design and synthesis of new herbicides and fungicides which will be eco-friendly. It PREFACE
Preface will aid the forensic scientist wishing to analyze a nondescript white powder is it heroin or flour? The list of scientific subject areas involving or its doval d go ing new mole on and on.Ore chemistr ting su bject since it leads to an esse tial nderstandingo The order in which the arly topics of this book are sented is i The first tw lust why doe structure and bor The third s ctional grou cr are to be capable ing e apparent actio whic 0r8a4 pounds on D nd E in by secti y, e b asic th eory on ms nucleophil ean e r ns can be u order.These look at the react the common functional groups which are important ir that studen ill find this textbook useful in their studies and ey have graspe what organic stry is all about they will read more widely and enter a truly exciting world of molecular science
will aid the forensic scientist wishing to analyze a nondescript white powder – is it heroin or flour? The list of scientific subject areas involving organic chemistry is endless – designing spacesuits, developing new photographic dyes, inventing new molecular technology in microelectronics – one could go on and on. Organic chemistry is an exciting subject since it leads to an essential understanding of molecules and their properties. The order in which the early topics of this book are presented is important. The first two sections cover structure and bonding, which are crucial to later sections. Just why does carbon form four bonds? What is hybridization? The third section on functional groups is equally crucial if students are to be capable of categorizing the apparent maze of reactions which organic compounds can undergo. It is followed by section D on stereochemistry, then sections E and F, in which the basic theory of reactions and mechanisms is covered. What are nucleophiles and electrophiles? What does a mechanism represent? What does a curly arrow mean? The remaining sections can be used in any order. These look at the reactions and mechanisms of the common functional groups which are important in chemistry and biochemistry. It is hoped that students will find this textbook useful in their studies and that once they have grasped what organic chemistry is all about they will read more widely and enter a truly exciting world of molecular science. x Preface
Section A-Structure and bonding A1 ATOMIC STRUCTURE OF CARBON Key Notes Atomic orbitals ,thesorbital shel orbitals are 2公金aee马 hape and can be assigned 2p2p,or 2p.depend- ing on the axis along which they are aligned. Energy levels The 1s orbital has a lower energy than the 2s orbital which has a lower energy than the 2p orbitals.The 2p orbitals have equal energy (i.e.they are degenerate). Electronic Carbon is in the second row of the periodic table and has six electrons which configuration will fill up lower energy atomic orbitals before entering higher energy orbitals(aufbau principle).Each orbital is allowed a maximum of two elec trons of opposite spin (Pauli exclusion principle).When orbitals of equal energy are available,electrons will occupy separate orbitals before pairing 2器2m.n，感he Related topic Covalent bonding and hybridization(A2) Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table.This means that there are two shells of atomic orbitals available for these electrons The first shell closest to the nucleus has a single s orbital-the 1s orbital.The second shell has a single s orbital (the 2s orbital)and three p orbitals (3 x 2p).Therefore,there are a total of five atomic orbitals into which these six electrons can fit thes orbitals an in shape with the 2s orbital being much large are dumbbell-shaped and are aligned along the x,yand z axes.Therefore,they are ssigned the and patomic orbitals (Fig.1). 4是@ Fig.1.Atomic orbitals
Section A – Structure and bonding A1 ATOMIC STRUCTURE OF CARBON Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table. This means that there are two shells of atomic orbitals available for these electrons. The first shell closest to the nucleus has a single s orbital – the 1s orbital. The second shell has a single s orbital (the 2s orbital) and three p orbitals (3 2p). Therefore, there are a total of five atomic orbitals into which these six electrons can fit. The s orbitals are spherical in shape with the 2s orbital being much larger then the 1s orbital. The p orbitals are dumbbell-shaped and are aligned along the x, y and z axes. Therefore, they are assigned the 2px, 2py and 2pz atomic orbitals (Fig. 1). Key Notes The atomic orbitals available for the six electrons of carbon are the s orbital in the first shell, the s orbital in the second shell and the three p orbitals in the second shell. The 1s and 2s orbitals are spherical in shape. The 2p orbitals are dumbbell in shape and can be assigned 2px, 2py or 2pz depending on the axis along which they are aligned. The 1s orbital has a lower energy than the 2s orbital which has a lower energy than the 2p orbitals. The 2p orbitals have equal energy (i.e. they are degenerate). Carbon is in the second row of the periodic table and has six electrons which will fill up lower energy atomic orbitals before entering higher energy orbitals (aufbau principle). Each orbital is allowed a maximum of two electrons of opposite spin (Pauli exclusion principle). When orbitals of equal energy are available, electrons will occupy separate orbitals before pairing up (Hund’s rule). Thus, the electronic configuration of a carbon atom is 1s 2 2s 2 2px 1 2py 1 . Related topic Covalent bonding and hybridization (A2) Electronic configuration Atomic orbitals Energy levels 1s 2s 2px 2py 2pz y z x y z x y z x y z x y z x Fig. 1. Atomic orbitals
2 Section A-Structure and bonding Energy levels The atomic orbitals described above are not of equal energy (Fig.2).The 1s orbital has the lowest energy.The 2s orbital is next in energy and the 2p orbitals have the highest energies.The three 2p orbitals have the same energy,meaning that they are degenerate. Energy 20 Fig.2.Energy levels of atomic orbitals. Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first.This is known as the aufbau princi- ple.The 1s orbital is filled up before the 2s orbital,which is filled up before the 2p orbitals.The Pauli exclusion principle states that each orbital is allowed a maxi- mum of two electrons and that these electrons must have fore the first four electrons fill un the 1s and 2s orbitals The electone rons in each and this i oppdown.There are tw nted in Fig.3 by drawing the a ns left to fit into the in: 2 e are two half-filled orbitals and one orbitals c 1 up once degenerate o wn Energy x↓2 ，—2p 1s Fig.3.Electronic configuration for carbon. The electronic configuration for carbon is 1s2s22p,2p,.The numbers in superscript refer to the numbers of electrons in each orbital.The letters refer to the types of atomic orbital involved and the numbers in front refer to which shell the orbital belongs
2 Section A – Structure and bonding Energy levels The atomic orbitals described above are not of equal energy (Fig. 2). The 1s orbital has the lowest energy. The 2s orbital is next in energy and the 2p orbitals have the highest energies. The three 2p orbitals have the same energy, meaning that they are degenerate. Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first. This is known as the aufbau principle. The 1s orbital is filled up before the 2s orbital, which is filled up before the 2p orbitals. The Pauli exclusion principle states that each orbital is allowed a maximum of two electrons and that these electrons must have opposite spins. Therefore, the first four electrons fill up the 1s and 2s orbitals. The electrons in each orbital have opposite spins and this is represented in Fig. 3 by drawing the arrows pointing up or down. There are two electrons left to fit into the remaining 2p orbitals. These go into separate orbitals such that there are two half-filled orbitals and one empty orbital. Whenever there are orbitals of equal energy, electrons will only start to pair up once all the degenerate orbitals are half filled. This is known as Hund’s rule. The electronic configuration for carbon is 1s 2 2s 2 2px 1 2py 1 . The numbers in superscript refer to the numbers of electrons in each orbital. The letters refer to the types of atomic orbital involved and the numbers in front refer to which shell the orbital belongs. Energy 1s 2s 2px 2py 2pz Fig. 2. Energy levels of atomic orbitals. Energy 1s 2s 2px 2py 2pz Fig. 3. Electronic configuration for carbon
Section A-Structure and bonding A2 COVALENT BONDING AND HYBRIDIZATION Key Notes Covalent When two hydrogen atoms approach each other,their Is atomic orbitals bonding interact to form a bonding and an antibonding molecular orbital(MO).A stable covalent bond is formed when the bonding MO is filled with a pair of electrons and the antibonding MO is empty. Sigma bonds Sigma (o)bonds are strong bonds with a circular cross-section formed by the head-on overlap of two atomic orbitals. Hybridization The electronic configuration of atomic carbon implies that carbon should form two bonds.However,it is known that carbon forms four bonds.When mix'the 2s and 2p orbitals of s hvbridiza ion.Ther Posibteypesofnbridaiom-ppadphybrid2atioh are three Related topics A) Covalent bonding Acovalent bond binds two atoms together in a molecula nic orbitals caularorbi one spe A atoms. om has a hal n two hydrogen the atoms the number of resulting MOs must equal the number of original atomic orbitals,Fig.1). Energy &gdngmoecularobtadl ⊙+ meamc = H:H ondingmolecular orbital Fig.1.Molecular orbitals for hydrogen (Ha). The MOsare of different enere One is more stable than the orinal atomic MO he c er is le an antibonding MO.The bondingM is shaped likea ruby ball and results from
Section A – Structure and bonding A2 COVALENT BONDING AND HYBRIDIZATION Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formed when atomic orbitals overlap to produce a molecular orbital – so called because the orbital belongs to the molecule as a whole rather than to one specific atom. A simple example is the formation of a hydrogen molecule (H2) from two hydrogen atoms. Each hydrogen atom has a half-filled 1s atomic orbital and when the atoms approach each other, the atomic orbitals interact to produce two MOs (the number of resulting MOs must equal the number of original atomic orbitals, Fig. 1). The MOs are of different energies. One is more stable than the original atomic orbitals and is called the bonding MO. The other is less stable and is called the antibonding MO. The bonding MO is shaped like a rugby ball and results from Key Notes When two hydrogen atoms approach each other, their 1s atomic orbitals interact to form a bonding and an antibonding molecular orbital (MO). A stable covalent bond is formed when the bonding MO is filled with a pair of electrons and the antibonding MO is empty. Sigma (σ) bonds are strong bonds with a circular cross-section formed by the head-on overlap of two atomic orbitals. The electronic configuration of atomic carbon implies that carbon should form two bonds. However, it is known that carbon forms four bonds. When carbon is part of an organic structure, it can ‘mix’ the 2s and 2p orbitals of the valence shell in a process known as hybridization. There are three possible types of hybridization – sp3 , sp2 and sp hybridization. Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3) sp2 Hybridization (A4) sp Hybridization (A5) Covalent bonding Sigma bonds Hybridization + 1s atomic orbital 1s atomic orbital = Bonding molecular orbital (full) Antibonding molecular orbital Energy (empty) H H H H HH HH Fig. 1. Molecular orbitals for hydrogen (H2 )