Introductory Bioelectronics

For Engineers and Physical Scientists

Ronald R. Pethig, Stewart Smith

• Publisher: Wiley
• Year: 2012
• ISBN:
  9781118443286
• Edition: 1
• Topic: Biological Sciences
• Subtopic: Biotechnology
• Index: Biological Sciences

Bioelectronics is a rich field of research involving the application of electronics engineering principles to biology, medicine, and the health sciences. With its interdisciplinary nature, bioelectronics spans state-of-the-art research at the interface between the life sciences, engineering and physical sciences.

Introductory Bioelectronics offers a concise overview of the field and teaches the fundamentals of biochemical, biophysical, electrical, and physiological concepts relevant to bioelectronics. It is the first book to bring together these various topics, and to explain the basic theory and practical applications at an introductory level.

The authors describe and contextualise the science by examining recent research and commercial applications. They also cover the design methods and forms of instrumentation that are required in the application of bioelectronics technology. The result is a unique book with the following key features:

• an interdisciplinary approach, which develops theory through practical examples and clinical applications, and delivers the necessary biological knowledge from an electronic engineer's perspective
• a problem section in each chapter that readers can use for self-assessment, with model answers given at the end of the book along with references to key scientific publications
• discussions of new developments in the bioelectronics and biosensors fields, such as microfluidic devices and nanotechnology

Supplying the tools to succeed, this text is the best resource for engineering and physical sciences students in bioelectronics, biomedical engineering and micro/nano-engineering. Not only that, it is also a resource for researchers without formal training in biology, who are entering PhD programmes or working on industrial projects in these areas.

Chapter 1

Basic Chemical and Biochemical Concepts

1.1 Chapter Overview

This chapter presents the background concepts of chemistry and thermodynamics of relevance to the subject of bioelectronics, and which are discussed further in most chapters of this book. The level of the material covered in this chapter is probably comparable to that covered by most students in pre-university basic chemistry courses. Graduates in engineering and the physical sciences may need to dig deeply into their recollections of such courses, and may also face new concepts. One objective here is to provide an awareness of some basic concepts of the chemical and energetic functioning of biological systems, of which even a modest understanding will go a long way to mastering the interdisciplinary field of bioelectronics.

After reading this chapter readers will gain a refreshed or new understanding of:

i. the formation of chemical bonds and how biological systems make use of the change in Gibbs free-energy ΔG of chemical reactions to perform the work required to retain their biological viability;

ii. chemical concentrations and activity coefficients;

iii. the concepts of nonpolar, polar, ionic, and hydrogen bonds;

iv. acids, bases and the biological importance of pH and buffers.

1.2 Energy and Chemical Reactions

1.2.1 Energy

A distinguishing characteristic of a living, rather than a nonliving, system is the ability to perform chemical transformations that produce fluxes of matter and energy. This process describes metabolism. Other characteristics that aid the identification of the living state are molecular organisation into systems of increasing complexity, and the abilities to self-produce and adapt to changes in environmental factors. The minimal level of organisation capable of exhibiting all these characteristics is the cell. The two principal forms of energy are kinetic and potential, associated with motion and stored energy, respectively. Kinetic energy in a molecular system can be interpreted in terms of the motions of its constituent molecules, which we term as heat. This heat can be determined indirectly by measuring the temperature of the molecular system. For heat to perform work (such as by an engine) it must flow from a region of higher to lower temperature. However, living systems are isothermal – they function at constant temperature and cannot utilise heat flow as a source of energy. Instead, living systems utilise the potential energy stored in the chemical bonds of molecules such as glucose or adenosine triphosphate (ATP). Cells continuously degrade such molecules, and the potential energy released when their chemical bonds are broken is used to perform various kinds of work, including the pumping of substances across membranes to produce chemical concentration gradients that in turn serve as sources of stored potential energy. This process, where chemical bond energy is converted into energy stored in the form of a chemical concentration gradient, is an example of the first law of thermodynamics which states that energy can neither be created nor destroyed. Other biological examples of this law include photosynthesis where the energy of sunlight absorbed by green leaves is converted into the chemical bond energy of glucose molecules, and in the conversion of chemical bond energy into mechanical and electrical energy by muscle cells and nerve cells, respectively.

All of the metabolic processes that produce the energy fluxes required for maintaining the living state involve the making and breaking of strong, covalent, chemical bonds between atoms in a molecule.

1.2.2 Covalent Chemical Bonds

Most biological molecules contain only six different atoms, namely carbon, hydrogen, oxygen, nitrogen, phosphorus and sulphur. The locations of these atoms in the Periodic Table of Elements are shown in Table 1.1. The electron shells of the atoms are labelled K, L and M. Each shell is composed of one or more subshells that represent the electronic orbitals about the nucleus of the atom. The first shell, K, has one subshell called the 1s shell and can accommodate a maximum of two electrons. The second shell, L, has two subshells (2s, 2p) that can accommodate a maximum of eight electrons, with six in the 2p shell. The third shell, M, has three subshells (3s, 3p, 3d) and can accommodate a maximum of 18 electrons, with 10 in the 3d shell.

Table 1.1 The locations of hydrogen (H), carbon (C), nitrogen (N), oxygen (O), phosphorus (P) and sulphur (S) in the Periodic Table of Elements.

Electrons in the outer shells have higher average energies than those in the inner shells, and their electron orbitals can extend farther from the nucleus. This contributes to how chemically reactive a particular atom may be in its interaction with other atoms. We can schematically represent the number and arrangement of electrons in the outer electron shells of these atoms as follows [1, 2]:

A covalent bond is formed by the sharing of unpaired electrons, one from the outer electron shell of each atom, between the nuclei of two atoms. These shared electrons then enter an electronic orbital that is common to both atoms, acting to reduce the repulsive force between the two positively charged nuclei and to hold them closely together. Thus, the hydrogen atom with one unpaired electron can form only one covalent bond, whilst carbon with four electrons forms four bonds. An example of this is methane (CH₄):

In the methane molecule the carbon atom is covalently bonded to four hydrogens.

In ethylene (C₂H₄) the two carbon atoms are held together by a double bond, and through the polymerisation of ethylene these double bonds are opened up to form the structure of polyethylene:

The nitrogen and phosphorus atoms possess five electrons in their outer electronic shells. These atoms can form either three covalent bonds (leaving a lone pair of unbonded electrons) or five covalent bonds. Examples include ammonia (NH₃) and phosphoric acid (H₃PO₄):

Oxygen contains six electrons in its outer electronic shell (known as the p-shell) and requires just two more electrons to completely fill this shell. It can accomplish this by forming two covalent bonds with another atom, such as in molecular oxygen (O₂) or in the carbonyl (C=O) chemical group:

The sulphur atom can also form two covalent bonds in this manner, as in hydrogen sulfide (H₂S). The outer electronic shell of the oxygen atom has two pairs of electrons that are not involved in covalent bond formation. This, however, does not apply to the sulphur atom, which can form as many as six covalent bonds as in sulphuric acid (H₂SO₄):

1.2.3 Chemical Concentrations

Concentrations of substances dissolved in solutions are often given in terms of weight/volume (e.g. mg/L, or mg/100 mL – a common clinical unit). These un...