File Name: introduction and scope of biochemistry .zip
Chemistry is the study of matter—what it consists of, what its properties are, and how it changes. Matter is anything that has mass and takes up space—that is, anything that is physically real. Some things are easily identified as matter—the screen on which you are reading this book, for example.
Roberts, J. Scott; Essential chemistry for biochemists. Essays Biochem 31 October ; 61 4 : — Within every living organism, countless reactions occur every second.
These reactions typically occur more rapidly and with greater efficiency than would be possible under the same conditions in the chemical laboratory, and while using only the subset of elements that are readily available in nature.
Despite these apparent differences between life and the laboratory, biological reactions are governed by the same rules as any other chemical reaction. Thus, a firm understanding of the fundamentals of chemistry is invaluable in biochemistry. There are entire textbooks devoted to the application of chemical principles in biological systems and so it is not possible to cover all of the relevant topics in depth in this short article. The aim is instead to provide a brief overview of those areas in chemistry that are most relevant to biochemistry.
We summarize the basic principles, give examples of how these principles are applied in biological systems and suggest further reading on individual topics. Biochemical systems carry out an enormous variety of chemical reactions with great efficiency. When a chemist carries out a reaction in the laboratory, they have numerous different techniques that can be applied to increase the yield or rate of a reaction; they can alter the temperature or pressure or perhaps add a catalyst.
In contrast, biological systems have to carry out the reaction at the temperature maintained by the organism, and, with the exception of organisms living deep in the ocean, at atmospheric pressure.
In biological systems, enzymes are employed to increase the rate of reaction; enzymes are proteins whose substrate-binding site acts to lower the energy of high energy species along the reaction pathway from starting material to product. They can achieve enormous enhancement in the rate of reaction compared with the uncatalysed reaction. The classic example is the enzyme triose phosphate isomerase which interconverts dihydroxyacetone phosphate and d- glyceraldehydephosphate during the breakdown of glucose.
This reaction occurs 10 9 times faster in the presence of the enzyme than when uncatalysed [ 1 ]. These rate enhancements are particularly remarkable when we consider that less than a third of the naturally occurring elements are used by biological systems. In order to be exploited in a biological system, elements must be sufficiently abundant in a form that can be taken up by living things. Thus, many catalytic species that are in common use in the laboratory are simply not available for biochemical reactions, for example palladium, used in the addition of hydrogen across a double bond.
Of the elements in the periodic table, 28 are essential for animal life Figure 1 ; the most recent element found to be essential is bromine, which was found to be required for the proper formation of networks by the protein collagen IV in [ 2 ]. Of the 28 essential elements, 11 make up The other 17 are known as trace elements and are present in very small amounts, ranging from a milligram to gram quantities in an adult human.
The 28 elements essential for animal life are indicated by coloured squares; trace elements are shown in yellow and those present in larger quantities are shown in green. Carbon, hydrogen, oxygen and nitrogen are the building blocks of organic biomolecules, calcium is present in large amounts in bones and teeth in addition to being vital for cell signalling in smaller amounts , phosphorus is likewise found in bones and teeth smaller quantities are a vital part of DNA, adenoside triphosphate — the energy currency of the cell — and play an important role in cell signalling.
Electrons in atoms are organized into a series of shells with successively higher energies and greater distance from the nucleus. The shells are identified by the principal quantum number which takes integer values from one to seven for the elements of the current periodic table. The shell with principle quantum number 1 has the lowest energy with 2 being the next highest in energy, and so on.
Within a shell there are subshells designated by the letters s, p, d and f, and within each subshell electrons occupy orbitals: regions of space that may be occupied by up to two electrons, and whose energy and shape can be described mathematically by an equation known as the wave function. It is important not to confuse orbitals with an orbit; electrons do not move around the nucleus along a fixed path as they would in an orbit. Instead the wave function allows us to calculate the probability of finding an electron at a particular position around the nucleus.
Different subshells have different numbers of orbitals; s subshells have just one orbital and can therefore accommodate only two electrons, p subshells have three orbitals which can be filled by six electrons in total, and d subshells have five orbitals accommodating up to ten electrons.
Shells are filled according to the Aufbau principle, i. Arrangements of electrons or electron configurations , in which the outer shell the occupied shell with highest principle quantum number is full are more favourable than those in which it is partially filled.
Of the elements in the periodic table, only the noble gases in group 18 have full outer shells. Bond formation, the movement of electrons between atoms, allows other elements to achieve this configuration. Chemical bonds can be covalent, where electrons are shared between atoms or ionic, where electrons are transferred from one atom to another resulting in one positively and one negatively charged species.
Looking at the number of electrons in the outer or valence shell enables us to work out how many bonds an atom would need to form in order to fill its outer shell.
It is important to note, however, that a bond will only actually form if the energy of the electrons in the bond is lower than the energy of those electrons in the isolated atoms. As the majority of biological chemistry relates to covalently bonded molecules composed primarily of the elements carbon, hydrogen, nitrogen and oxygen, it is particularly important to know how many bonds each of these elements form.
Hydrogen, with its single valence electron requires one additional electron to achieve the noble gas configuration and therefore makes only one bond. Carbon, with four valence electrons, achieves a full outer shell by forming four covalent bonds, for example by sharing an electron with each of four other atoms. Nitrogen has five valence electrons and forms three covalent bonds leaving one pair of non-bonded electrons; a lone pair.
The lone pair is important for the reactivity of nitrogen as it can be used to make a new bond with electron-poor species in chemical reactions.
Finally, oxygen with six electrons makes two covalent bonds and has two lone pairs of electrons. When an ionic bond is formed, electrons are transferred completely from one atom to the other. The interaction in an ionic bond is entirely Coulombic in nature i. Such a bond occurs when the elements differ widely in their ability to attract electrons or more formally, when there is a very large energy difference between the valence orbitals the outermost orbitals containing the electrons available to participate in bonding in the two atoms.
Covalent bonding involves sharing of electrons between atoms and occurs when the two atoms are more similar in their ability to attract electrons; i. Sharing of electrons in a covalent bond requires atomic orbitals on each of the atoms to interact with each other. One of the consequences of this is that, in contrast with ionic bonds, covalent bonds are directional: a new bond cannot form in a region of space where the orbitals that interact to form the new bond have no electron density.
Although a pair of electrons are shared when a covalent bond is formed, however, this sharing is not necessarily equal. Atoms with a strong ability to attract electrons, i. If the two atoms forming a covalent bond differ in electronegativity then there will be greater electron density closer to the more electronegative atom. This results in a permanent dipole, with one atom partially negatively charged and the other partially positively charged Figure 2 C. In these images, atomic orbitals and bonds are depicted as line drawings indicating shape and as isosurfaces, regions of space enclosing a defined fraction of the electron density.
The description of covalent bonding has so far assumed that a covalent bond involves the sharing of one pair of electrons between two atoms, i. For the majority of biological molecules this description is adequate, however in some cases this description of bonding does not explain the observed properties of a molecule.
A well-known example is benzene. It was originally thought that benzene contained three alternating single and double bonds, however measurements showed that the bonds were all of equal length.
We now consider the bonding in benzene not as three pairs of p-orbitals each interacting to make one double bond, but six p-orbitals each interacting with its neighbours to create a ring of electron density above and below the plane of the carbon atoms Figure 3 A. Compounds with delocalized rings of electrons are of major importance in biological systems. A Each of the six carbon atoms in benzene contribute a p-orbital to the delocalized system.
A circle in the centre of the ring can be used to highlight the fact that the system is delocalized, however many biochemists prefer to use the alternating bond representation. B The four bases found in DNA, all have delocalized rings of p-orbitals; in this diagram atoms shown in red each contribute a p-orbital to the delocalized system.
It is not immediately obvious that the atoms indicated with an arrow can contribute a p-orbital, however using a more sophisticated approach to bonding we can show that this is the case. Adenine and guanine each have ten electrons in a delocalized ring, while cytosine and thymine have six. C Representation of the delocalized carboxylate anion.
In this system, four electrons two from the double bond and two from the negatively charged oxygen are delocalized over three atoms. D Retinal has a linear delocalized system including 12 p-orbitals. Each of the atoms that contributes a p-orbital to the delocalized system is shown in red.
Delocalization of electrons does not only occur in rings; another type of system where delocalization occurs is where three or more parallel p-orbitals are adjacent. Consider the carboxylate anion discussed in The carbonyl functional group section where a carbon atom makes a double bond with one oxygen atom and a single bond with a negatively charged oxygen atom. In this structure, we can visualize the negative charge on the oxygen being used to make a new double bond with the carbon atom and the existing double bond breaking to leave a negative charge on the oxygen atom Figure 3 C.
Although we can visualize single and double bonds exchanging position within the molecule, this is not an accurate description of bonding in the molecule. In reality, the electron density is spread over three p-orbitals, and a higher electron density exists on the two oxygen atoms than on the central carbon atom. Delocalization of electron density across three p-orbitals is also important in explaining why the bond formed between two amino acids in a protein chain is planar see Functional groups found in amino acids section.
Molecules with electrons are delocalized over a larger number of adjacent parallel p-orbitals are also common in biology. These molecules are usually referred to as conjugated and can be identified by their alternating chain of single and double bonds.
For example, retinal, the light-absorbing molecule that is bound to the protein opsin in the photoreceptor cells responsible for vision in mammals, has electron density delocalized across 12 p-orbitals Figure 3 D. The long delocalized system is essential for the absorption of light in the visible region of the electromagnetic spectrum. These interactions are much weaker than the covalent bond but they occur very frequently and, as a result, can have a huge influence on the properties of a molecule.
Many biomolecules are macromolecules with thousands of atoms and therefore make many hundreds of thousands of non-covalent interactions. Non-covalent interactions are particularly important in proteins. Proteins are polymers of amino acids, synthesized in a linear chain on the ribosome. Each protein chain folds into a specific 3D structure that is essential for its function; non-covalent interactions between the constituent amino acids determine the 3D structure.
Non-covalent interactions are also important in DNA where they help to ensure that the sequence of DNA is preserved upon replication; in the lipid bilayer where non-covalent interactions between lipids create a barrier around the cell; and in molecular recognition discussed in more detail in [ 3 ].
There are several classes of non-covalent interactions; here we discuss van der Waals interactions, hydrogen bonds and briefly, ionic interactions.
A dipole is an uneven distribution of electron density within a molecule such that one region of the molecule has a higher electron density than the other and the two regions are equally but oppositely charged. These dipoles can be permanent or instantaneous. Permanent dipoles occur due to the uneven charge distribution in a covalent bond between two elements that differ greatly in electronegativity. Interactions between instantaneous dipoles are called London dispersion forces. They are the weakest among the non-covalent interactions, but also the most prevalent.
London dispersion forces occur because the electron density in an atom or molecule does not have an even distribution; at any one time the electron density may be higher in one region than the other. The electron density is redistributed with time, thus the regions of high electron density are different from one moment to the next. The uneven charge distribution is called instantaneous dipole. The distribution of the electron density in neighbouring molecules is influenced by the dipole of the first molecule; areas of relative high electron density on one molecule induce an area of low electron density on the neighbouring molecule and vice versa; thus neighbouring molecules form instantaneous dipoles that attract each other.
When molecules have large surface areas that can come into close contact, for example in biological macromolecules, these interactions can make a huge contribution to the total free energy. Hydrogen bonds are a special case of dipole—dipole interaction, but are considered separately here as they are vital for the function of many biochemical systems.
Roberts, J. Scott; Essential chemistry for biochemists. Essays Biochem 31 October ; 61 4 : — Within every living organism, countless reactions occur every second. These reactions typically occur more rapidly and with greater efficiency than would be possible under the same conditions in the chemical laboratory, and while using only the subset of elements that are readily available in nature. Despite these apparent differences between life and the laboratory, biological reactions are governed by the same rules as any other chemical reaction.
Biochemistry is the application of chemistry to the study of biological processes at the cellular and molecular level. It emerged as a distinct discipline around the beginning of the 20th century when scientists combined chemistry, physiology, and biology to investigate the chemistry of living systems. Biochemistry is both life science and a chemical science - it explores the chemistry of living organisms and the molecular basis for the changes occurring in living cells. It uses the methods of chemistry,. It has provided explanations for the causes of many diseases in humans, animals and plants.
Biochemistry or biological chemistry , is the study of chemical processes within and relating to living organisms. Over the last decades of the 20th century, biochemistry has become successful at explaining living processes through these three disciplines. Almost all areas of the life sciences are being uncovered and developed through biochemical methodology and research. Much of biochemistry deals with the structures, functions, and interactions of biological macromolecules , such as proteins , nucleic acids , carbohydrates , and lipids. They provide the structure of cells and perform many of the functions associated with life.
The scope of biochemistry is as wide as life itself. Wherever there is life, chemical processes are occurring. Biochemists study the chemical processes that occur in.
Biochemistry , study of the chemical substances and processes that occur in plants , animals , and microorganisms and of the changes they undergo during development and life. It deals with the chemistry of life, and as such it draws on the techniques of analytical , organic, and physical chemistry , as well as those of physiologists concerned with the molecular basis of vital processes. All chemical changes within the organism—either the degradation of substances, generally to gain necessary energy, or the buildup of complex molecules necessary for life processes—are collectively called metabolism.
Biochemistry emerged as a separate discipline when scientists combined biology with organic, inorganic, or physical chemistry and began to study such topics as how living things obtain energy from food, the chemical basis of heredity, and what fundamental changes occur in disease. Biochemistry includes the sciences of molecular biology; immunochemistry; neurochemistry; and bioinorganic, bioorganic, and biophysical chemistry. Biochemistry, sometimes called biological chemistry, is the study of chemical processes in living organisms, including, but not limited to, living matter. Biochemistry governs all living organisms and living processes.
Biochemistry in broad terms is the study of the chemical composition of the living matter and the biochemical processes that underlie life activities during growth and maintenance. This article is an attempt to explore the metamorphosis of biochemistry from a pupa entwined in its own cocoon to a vibrantly colored phenomenon. Studies pertaining to this discipline of science began with Biochemistry interfaces with biology and chemistry even before nineteenth century with studies concerned with the chemical processes that take place within living cells. Modern biochemistry developed out of and largely came to replace what in the nineteenth and early twentieth centuries was called physiological chemistry, which dealt more with extra cellular chemistry, such as the chemistry of digestion and of body fluids.
Physiology is the study of normal function within living creatures. It is a sub-section of biology, covering a range of topics that include organs, anatomy, cells, biological compounds, and how they all interact to make life possible. From ancient theories to molecular laboratory techniques, physiological research has shaped our understanding of the components of our body, how they communicate, and how they keep us alive. Here are some key points about physiology.
Photo: a Biochemistry Lab. Biochemistry, sometimes called biological chemistry is the study of the structure, composition, and chemical reactions of substances in living systems. By controlling information flow through biochemical signaling and the flow of chemical energy through metabolism, biochemical processes give rise to the complexity of life. Biochemistry emerged as a separate discipline when scientists combined biology with organic, inorganic, and physical chemistry and began to study how living things obtain energy from food, the chemical basis of heredity, what fundamental changes occur in disease, and related issues. Biochemistry includes the sciences of molecular biology, immunochemistry, and neurochemistry, as well as bioinorganic, bioorganic, and biophysical chemistry. But today, the main focus of pure biochemistry is on understanding how biological molecules give rise to the processes that occur within living cells, which in turn relates greatly to the study and understanding of tissues, organs, and whole organisms—that is, all of biology. Introduction and Definition of Biochemistry The most conspicuous attribute o living organisms is that they are complicated and highly organized.
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