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- Herausgeber: John Wiley & Sons
- Kategorie: Wissenschaft und neue Technologien
- Sprache: Englisch
Essential Microbiology 2nd Edition is a fully revised comprehensive introductory text aimed at students taking a first course in the subject. It provides an ideal entry into the world of microorganisms, considering all aspects of their biology (structure, metabolism, genetics), and illustrates the remarkable diversity of microbial life by devoting a chapter to each of the main taxonomic groupings. The second part of the book introduces the reader to aspects of applied microbiology, exploring the involvement of microorganisms in areas as diverse as food and drink production, genetic engineering, global recycling systems and infectious disease.
Essential Microbiology explains the key points of each topic but avoids overburdening the student with unnecessary detail. Now in full colour it makes extensive use of clear line diagrams to clarify sometimes difficult concepts or mechanisms. A companion web site includes further material including MCQs, enabling the student to assess their understanding of the main concepts that have been covered.
This edition has been fully revised and updated to reflect the developments that have occurred in recent years and includes a completely new section devoted to medical microbiology. Students of any life science degree course will find this a concise and valuable introduction to microbiology.
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Veröffentlichungsjahr: 2013
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Contents
Cover
Title Page
Copyright
Preface to Second Edition
Preface to First Edition
Acknowledgements
About the Companion Website
Part I: Introduction
Chapter 1: Microbiology: What, Why and How?
Chapter 2: Biochemical Principles
Chapter 3: Cell Structure and Organisation
Part II: Microbial Nutrition, Growth and Metabolism
Chapter 4: Microbial Nutrition and Cultivation
Chapter 5: Microbial Growth
Chapter 6: Microbial Metabolism
Part III: Microbial Diversity
Chapter 7: Prokaryote Diversity
Chapter 8: The Fungi
Chapter 9: The Protista
Chapter 10: Viruses
Part IV: Microbial Genetics
Chapter 11: Microbial Genetics
Chapter 12: Microorganisms in Genetic Engineering
Part V: Microorganisms in the Environment
Chapter 13: Microbial Associations
Chapter 14: Microorganisms in the Environment
Part VI: Medical Microbiology
Chapter 15: Human Microbial Diseases
Chapter 16: The Control of Microorganisms
Chapter 17: Antimicrobial Agents
Part VII: Microorganisms in Industry
Chapter 18: Industrial and Food Microbiology
Glossary
Further Reading
Index
This edition first published 2013 © 2013 by John Wiley & Sons, Ltd
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Library of Congress Cataloging-in-Publication Data
Hogg, Stuart (Stuart I.)Essential microbiology / Stuart Hogg. – 2nd ed.p. cm.Includes index.ISBN 978-1-119-97891-6 (cloth) – ISBN 978-1-119-97890-9 (pbk.)I. Title.[DNLM: 1. Microbiological Phenomena. 2. Microbiological Techniques. 3. Microbiology.QW 4]
579–dc23 2012051595
A catalogue record for this book is available from the British Library.
Wiley also publishes its books in a variety of electronic formats. Some content that appears in print may not be available in electronic books.
Cover image: Gary Thompson Cover design by Stain 2.png supplied by publisher
Preface to Second Edition
It is now seven years since the first edition of Essential Microbiology was published, so it is high time the contents were updated, and I have taken the opportunity to revise the layout in the hope that it will better serve its target readership. The main change to the book from its original incarnation is the inclusion of a chapter on microbial disease in humans. When preparing the content of the first edition, the one major area of doubt I had was whether or not to include a chapter or section on medical microbiology. I was urged to do so by a number of colleagues, but in the end I resisted, feeling it to be too large a topic for inclusion in a general introductory text. The invitation to prepare a second edition, however, has given me an opportunity to reconsider the matter, and comments from several reviewers, together with further reflection on my own part, have persuaded me to change my mind. I have therefore introduced a new chapter on microbial disease in humans, supplementing new material with some expanded and repackaged from other chapters in the first edition. This has resulted in a shuffling and reordering of the second half of the book, which I hope leads to a more logical structure. The new edition no longer features end-of-chapter quizzes; however, these and other forms of self-assessment can now be found on the book's dedicated website. The other major change that will be noticed by anyone familiar with the original book is the introduction of colour. I feel strongly that a book such as this should be visually attractive as well as instructive, and am grateful to my editorial team at Wiley for allowing me this indulgence, in spite of the additional pressure it creates in trying to keep the selling price to a minimum, which was always one of the principal aims of Essential Microbiology.
As always, I should be grateful to receive any comments and suggestions for improvement from students or their tutors.
Stuart HoggSeptember 2012
Preface to First Edition
Every year, in UK universities alone, many hundreds of students study microbiology as part of an undergraduate course. For some, the subject will form the major part of their studies, leading to a BSc degree in Microbiology, or a related subject such as Bacteriology or Biotechnology. For the majority, however, the study of microbiology will be a brief encounter, forming only a minor part of their course content.
A number of excellent and well-established textbooks are available to support the study of microbiology; such titles are mostly over 1000 pages in length, beautifully illustrated in colour, and rather expensive. This book in no way seeks to replace or compete with such texts, which will serve specialist students well throughout their three years of study, and represent a sound investment. It is directed rather towards the second group of students, who require a text that is less detailed, less comprehensive, and less expensive! The majority of the students in my own classes are enrolled on BSc degrees in Biology, Human Biology and Forensic Science; I have felt increasingly uncomfortable about recommending that they invest a substantial sum of money on a book much of whose content is irrelevant to their needs. Alternative recommendations, however, are not thick on the ground. This, then, was my initial stimulus to write a book of ‘microbiology for the non-microbiologist’.
The facts and principles you will find here are no different from those described elsewhere, but I have tried to select those topics that one might expect to encounter in years 1 and 2 of a typical non-specialist degree in the life sciences or related disciplines. Above all, I have tried to explain concepts or mechanisms; one thing my research for this book has taught me is that textbooks are not always right, and they certainly don't always explain things as clearly as they might. It is my wish that the present text will give the attentive reader a clear understanding of sometimes complex issues, whilst avoiding over-simplification.
The book is arranged into seven sections, the fourth of which, Microbial Genetics, acts as a pivot, leading from principles to applications of microbiology. Depending on their starting knowledge, readers may ‘dip into’ the book at specific topics, but those whose biological and chemical knowledge is limited are strongly recommended to read Chapters 2 and 3 for the foundation necessary for the understanding of later chapters. Occasional boxes are inserted into the text, which provide some further enlightenment on the topic being discussed, or offer supplementary information for the inquisitive reader. As far as possible, diagrams are limited to simple line drawings, most of which could be memorised for reproduction in an examination setting. Although a Glossary is provided at the end of the book, new words are also defined in the text at the point of their first introduction, to facilitate uninterrupted reading. All chapters except the first are followed by a self-test section in which readers may review their knowledge and understanding by ‘filling in the gaps’ in incomplete sentences; the answers are all to be found in the text, and so are not provided separately. The only exceptions to this are two numerical questions, the solutions to which are to be found at the back of the book. By completing the self-test questions, the reader effectively provides a summary for the chapter.
A book such as this stands or falls by the reception it receives from its target readership. I should be pleased to receive any comments on the content and style of Essential Microbiology from students and their tutors, all of which will be given serious consideration for inclusion in any further editions.
Stuart HoggJanuary 2005
Acknowledgements
I would like to thank those colleagues who took the time to read over individual chapters of this book, and those who reviewed the entire manuscript. Their comments have been gratefully received, and in some cases spared me from the embarrassment of seeing my mistakes perpetuated in print.
Thanks are also due to my editorial team at John Wiley, Rachael Ballard and Fiona Seymour, and production editor Jasmine Chang for ensuring smooth production of this book.
I am grateful to those publishers and individuals who have granted permission to reproduce diagrams. Every effort has been made to trace holders of copyright; any inadvertent omissions will gladly be rectified in any future editions of this book.
Finally, I would like to express my gratitude to my family for allowing me to devote so many weekends to ‘the book’.
About the Companion Website
This book is accompanied by a companion website:
www.wiley.com/go/hogg/essentialmicrobiology
The website includes:
Powerpoints of all figures from the book for downloadingPDFs of tables from the bookShort Answer QuestionsKey conceptsLinks for further readingI
Introduction
2
Biochemical Principles
All matter, whether living or non-living, is made up of atoms; the atom is the smallest unit of matter capable of entering into a chemical reaction. Atoms can combine together by bonding, to form molecules, which range from the small and simple to the large and complex. The latter are known as macromolecules; major cellular constituents such as carbohydrates and proteins belong to this group, and it is with these that this chapter is mainly concerned (Table 2.1). In order to appreciate how these macromolecules operate in the structure and function of microbial cells, however, you need to appreciate the basic principles of how atoms are constructed and how they interact with one other.
Table 2.1Biological macromolecules. Important examples of each of the four major classes of macromolecule found in biological systems
2.1 Atomic structure
All atoms have a central, positively charged nucleus, which is very dense and makes up most of the mass of the atom. The nucleus is made up of two types of particle, protons and neutrons. Protons carry a positive charge, and neutrons are uncharged, hence the nucleus overall is positively charged. It is surrounded by much lighter, and rapidly orbiting, electrons (Figure 2.1). These are negatively charged, with the charge on each electron being equal (but of course opposite) to that of the protons; however, the electrons have only 1/1840 of the mass of either protons or neutrons. It is the attractive force between the positively charged protons and the negatively charged electrons that holds the atom together.
Figure 2.1Atomic structure. The nucleus of a carbon atom contains six protons and six neutrons, surrounded by six electrons. Note how these are distributed between inner (2) and outer (4) electron shells.
The number of protons in the nucleus is called the atomic number, and ranges from one to over one hundred. The combined total of protons and neutrons is known as the mass number. All atoms have an equal number of protons and electrons, so regardless of the atomic number, the overall charge on the atom will always be zero.
Atoms having the same atomic number have the same chemical properties; such atoms all belong to the same element. An element is made up of one type of atom only and cannot be chemically broken down into simpler substances; thus pure copper, for example, is made up entirely of copper atoms. There are 92 of these elements, 26 of which commonly occur in living things. Each element has been given a universally agreed symbol; examples that we shall encounter in biological macromolecules include carbon (C), hydrogen (H) and oxygen (O). The atomic numbers of selected elements are shown in Table 2.2.
Table 2.2Symbols and atomic numbers of some elements occurring in living systems
ElementSymbolAtomic no.Hydrogen H 1Carbon C 6Nitrogen N 7Oxygen O 8Sodium Na11Magnesium Mg12Phosphorus P15Sulphur S16Chlorine Cl17Potassium K19Iron Fe26The relationship between neutrons, protons, atomic number and mass number is illustrated in Table 2.3. We have used carbon as an example, since all living matter is based upon this element. The carbon represented can be expressed in the form:
Table 2.3The vital statistics of carbon
The number of neutrons in an atom can be deduced by subtracting the atomic number from the mass number. In the case of carbon, this is the same as the number of protons (6), but this is not always so; phosphorus for example has 15 protons and 16 neutrons, giving it an atomic number of 15 and a mass number of 31.
2.1.1 Isotopes
Although the number of protons in the nucleus of a given element is always the same, the number of neutrons can vary, giving different forms, or isotopes, of that element. Carbon-14 (14C) is a naturally occurring but rare isotope of carbon that has eight neutrons instead of six, hence the atomic mass of 14. Carbon-13 (13C) is a rather more common isotope, making up around 1% of naturally occurring carbon; it has seven neutrons per atomic nucleus. The atomic mass (or atomic weight) of an element is the average of the mass numbers of an element's different isotopes, taking into account the proportions in which they occur. Carbon-12 is by far the predominant form of the element in nature, but the existence of small amounts of the other (slightly heavier) forms means that the atomic mass is 12.011 rather than exactly 12. Some isotopes are stable, while others decay spontaneously, with the release of subatomic particles. The latter are called radioisotopes; 14C is a radioisotope, while the other two forms of carbon are stable isotopes. Radioisotopes have been an extremely useful research tool in a number of areas of molecular biology.
The electrons that orbit around the nucleus do not do so randomly, but are arranged in a series of electron shells, radiating out from the nucleus (see Figure 2.1). These layers correspond to different energy levels, with the highest energy levels being located furthest away from the nucleus. Each shell can accommodate a maximum number of electrons, and they always fill up starting at the innermost one, that is, the one with the lowest energy level. In our example, carbon has filled the first shell with two electrons, and occupied four of the eight available spaces on the second.
The chemical properties of atoms are determined by the number of electrons in the outermost occupied shell. Neon, one of the so-called ‘noble’ gases, has an atomic number of 10, completely filling the first two shells, and is chemically unreactive or inert. Atoms that achieve a similar configuration are all stable. If, on the other hand, such an arrangement is not achieved, the atom is unstable, or reactive. Reactions take place between atoms that attempt to achieve stability by attaining a full outer shell. These reactions may involve atoms of the same element or ones of different elements; the result in either case is a molecule or ion. Figure 2.2 shows how atoms combine to form a molecule. A substance made up of molecules containing two or more different elements is called a compound. In each example, the product of the reaction has a full outer electron shell; note that some atoms are donating electrons, while others are accepting them.
Figure 2.2Atoms join to form molecules. The formation of (a) hydrogen and (b) methane by covalent bonding. Each atom achieves a full set of electrons in its outer shell by sharing with another atom. A shared pair of electrons constitutes a covalent bond.
If most of the spaces in the outermost electron shell are full, or if most are empty, atoms tend to strive for stability by gaining or losing electrons, as shown in Figure 2.3. When this happens, an ion is formed, which carries either a positive or negative charge. Positively charged ions are called cations and negatively charged ones anions. The sodium atom, for example, has 11 electrons, meaning that the inner two electron shells are filled and a lone electron occupies the third shell. When it loses this last electron, it has more protons than electrons, and therefore has a net positive charge of one; when this happens, it becomes a sodium ion, Na+ (see Figure 2.3).
Figure 2.3Ion formation. Sodium achieves stability by losing the lone electron from its outermost shell. The resulting sodium ion Na+ has 11 protons and 10 electrons, hence it carries a single positive charge. Chlorine becomes ionised to chloride (Cl−) when it gains an electron to complete its outer shell.
2.1.2 Chemical bonds
The force that causes two or more atoms to join together is known as a chemical bond, and several types are found in biological systems. The interaction between sodium and chloride ions shown in Figure 2.4 is an example of ionic bonding, where the transfer of an electron from one party to another means that both achieve a complete outer electron shell. There is an attractive force between positively and negatively charged ions, called an ionic bond. Certain elements form ions with more than a single charge, by gaining or losing two or more electrons in order to achieve a full outer electron shell; thus calcium ions (Ca2+) are formed by the loss of two electrons from a calcium atom.
Figure 2.4Ionic bonding. A positively charged Na+ and negatively charged Cl− attract each other, and an ionic bond is formed. The result is a molecule of sodium chloride.
The goal of stability through a full complement of outer shell electrons may also be achieved by means of sharing one or more pairs of electrons. Consider the formation of methane (see Figure 2.2); a carbon atom, which has four spaces in its outer shell, can achieve a full complement by sharing electrons from four separate hydrogen atoms. This type of bond is a covalent bond.
Sometimes, a pair of atoms share not one but two pairs of electrons (Figure 2.5). This involves the formation of a double bond. Triple bonding, through the sharing of three pairs of electrons, is also possible, but rare.
Figure 2.5Double bond formation. In the formation of carbon dioxide, the carbon atom shares two pairs of electrons with each oxygen atom.
In the examples of covalent bonding we've looked at so far, the sharing of the electrons has been equal, but this is not always the case because sometimes the electrons may be drawn closer to one atom than another (Figure 2.6a). This has the effect of making one atom slightly negative and another slightly positive. Molecules like this are called polar molecules and the bonds are polar bonds. Sometimes a large molecule may have both polar and nonpolar areas. Polar molecules are attracted to each other, with the negative areas of one molecule attracted to the positive areas of another (Figure 2.6b). In water, hydrogen atoms bearing a positive charge are drawn to the negatively charged oxygens.
Figure 2.6Water is a polar molecule. (a) The electrons of the hydrogen atoms are strongly attracted to the oxygen atom, causing this part of the water molecule to carry a slightly negative charge, and the hydrogen part a slightly positive one. (b) Because of their polar nature, water molecules are attracted to each other by hydrogen bonding. Hydrogen bonding is much weaker than ionic or covalent bonding, but plays an important role in the structure of macromolecules such as proteins and nucleic acids.
This attraction between polar atoms is called hydrogen bonding, and can take place between covalently bonded hydrogen and any electronegative atom, most commonly oxygen or nitrogen. Hydrogen bonds are much weaker than either ionic or covalent bonds; however, if sufficient of them form in a compound, the overall bonding force can be appreciable. Each water molecule can form hydrogen bonds with others of its kind in four places. In order to break all these bonds, a substantial input of energy is required, explaining why water has such a relatively high boiling point, and why most of the water on our planet is in liquid form.
Another weak form of interaction is brought about by van der Waals forces, which occur briefly when two nonpolar molecules (or parts of molecules) come into very close contact with one another. Although transient, and generally even weaker than hydrogen bonds, they occur in great numbers in certain macromolecules and play an important role in holding proteins together (see Section 2.3.3).
Water is essential for living things, both in the composition of their cells and in the environment surrounding them. Organisms are made up of 60–95% water by weight, and even inert dormant forms like spores and seeds have a significant water component. This dependence on water is a function of its unique properties, which in turn derive from its polar nature.
Water is the medium in which most biochemical reactions take place; it is a highly efficient solvent; indeed, more substances will dissolve in water than in any other solvent. Substances held together by ionic bonds tend to dissociate into anions and cations in water, because as individual solute molecules become surrounded by molecules of water, hydration shells are formed, in which the negatively charged parts of the solute attract the positive region of the water molecule, and the positive parts the negative region (Figure 2.7). The attractive forces that allow the solute to dissolve are called hydrophilic forces, and substances that are water-soluble are hydrophilic (water-loving). Other polar substances such as sugars and proteins are also soluble in water by forming hydrophilic interactions.
Figure 2.7Ions dissolve in water. An ionic compound such as sodium chloride dissociates in water to its constituent ions. Water molecules form hydration shells around both Na+ and Cl− ions.
Molecules such as oils and fats are nonpolar, and because of their non-reactivity with water are termed hydrophobic (‘water-fearing’). If such a molecule is mixed with water, it will be excluded, as water molecules ‘stick together’. This very exclusion by water can act as a cohesive force among hydrophobic molecules (or hydrophobic areas of large molecules). This is often called hydrophobic bonding, but it isn't really bonding as such, rather a shared avoidance of water. All living cells have a hydrophilic interior surrounded by a hydrophobic membrane, as we'll see in Chapter 3.
An amphipathic substance is one that is part polar and part nonpolar. When such a substance is mixed with water, micelles are formed (Figure 2.8); the nonpolar parts are excluded by the water and group together as described above, leaving the polar groups pointing outwards into the water, where they are attracted by hydrophilic forces. Detergents exert their action by trapping insoluble grease inside the centre of a micelle, while interaction with water allows them to be rinsed away (see Chapter 16).
Figure 2.8Amphipathic molecules form micelles in water. In an aqueous environment, amphipathic substances align their molecules so that the nonpolar parts are hidden away from the water. Reproduced from Black, JG (1999) Microbiology: Principles and Explorations, 4th edn, with permission from John Wiley & Sons.
Water plays a role in many essential metabolic reactions, and its polar nature allows for the breakdown to hydrogen and hydroxyl ions (H+ and OH−), and resynthesis as water. Water acts as a reactant in hydrolysis reactions such as:
or as a product in certain synthetic reactions, like
2.2 Acids, bases and pH
Only a minute proportion of water molecules, something like one in every 500 million, are present in the dissociated form, but as we've already seen, the H+ and OH− ions play an important part in cellular reactions.
A solution becomes acid or alkaline if there is an imbalance in the amount of these ions present. If there is an excess of H+, the solution becomes acid, whilst if OH− predominates, it becomes alkaline. The pH of a solution is an expression of the molar concentration of hydrogen ions; it is expressed thus:
In pure water, hydrogen ions are present at a concentration of 10−7 M, thus the pH is 7.0. This is called neutrality, where the solution is neither acid nor alkaline. At higher concentrations of H+, such as 10−3 M (1 millimolar), the pH value is lower, in this case 3.0, so acid solutions have a value below 7. Conversely, alkaline solutions have a pH above 7. You will see from this example that an increase of 104 (10 000)-fold in the [H+] leads to a change of only four points on the pH scale. This is because it is a logarithmic scale, thus a solution of pH 10 is 10 times more alkaline than one of pH 9, and 100 times more than one of pH 8. Figure 2.9 shows the pH value of a number of familiar substances.
Figure 2.9The pH value of some common substances. Most microorganisms exist at pH values around neutrality, but representatives are found at extremes of both acidity and alkalinity. Reproduced from Black, JG (1999) Microbiology: Principles and Explorations, 4th edn, with permission from John Wiley & Sons.
Most microorganisms live in an aqueous environment, and the pH of this is very important. Most will only tolerate a small range of pH, and the majority occupy a range around neutrality, although as we shall see later in this book, there are some startling exceptions to this.
Most of the important molecules involved in the chemistry of living cells are organic, that is, they are based on a skeleton of covalently linked carbon atoms. Biological molecules have one or more functional groups
