Biochemical concepts
3.1 Organization of matter
3.1.1 Matter and elements
All things, living and non-living, consist of mat- ter, defined as anything that occupies space and has mass. Matter, in turn, is made up of chemical building blocks known as elements – substances which cannot be split into simpler substances by ordinary chemical means. There are 112 known chemical elements, 92 of which occur naturally on Earth. Of these 92 elements, there are 26 which occur normally in our bodies. Each element can be symbolized using a one- or two-letter abbreviation which, in the following paragraphs, we include in brackets after stating the name of the element in question.
The most abundant elements which make up our bodies (constituting approximately 96% of our body’s mass) areoxygen(O),carbon(C),hydro- gen (H) andnitrogen(N).
Eight other elements make up 3.8% of our body’s mass. These include calcium (Ca), phosphorous (P), potassium (K), sulphur (S), sodium (Na), chlorine (Cl), magnesium (Mg) and iron (Fe).
The remaining 0.8% of the body’s store of ele- ments consists of 14 elements known as the trace elements. These include aluminium (Al), boron (B), chromium (Cr), cobalt (Co), copper (Cu), fluo- rine (F), iodine (I), manganese (Mn), molybdenum (Mo), selenium (Se), silicon (Si), tin (Sn), vana- dium (V) and zinc (Zn).
Elements are important for human life. They are components of the energy sources of car- bohydrates, fats and proteins, and they also provide us withwaterand importantvitaminsand mineralsthat are needed to sustain healthy living.
An overview of the main chemical elements within the human body is shown in Table 3.1.
3.1.2 Atoms and atomic structure
Each element is made from atoms, which are the smallest units of matter. Atoms are extremely
Table 3.1 An overview of the body’s main chemical elements and some of their known functions
Chemical element Percentage abundance Function
(symbol) in the human body
Oxygen (O) 65 Used to generate energy using aerobic processes and also part of water and energy sources such as carbohydrates, fats and proteins.
Carbon (C) 18.5 Important component of organic (i.e. carbon- containing) molecules such as carbohydrates, fats, proteins and deoxyribonucleic acids (DNA – a cell’s genetic material).
Hydrogen (H) 9.5 Component of water and most organic molecules.
Nitrogen (N) 3.2 Component of all proteins and DNA.
Calcium (Ca) 1.5 Important component for maintenance of bones and teeth; also involved in regulatory processes such as hormone release, muscle contraction and enzyme activation.
Phosphorous (P) 1 Component of ATP (the immediate energy supply for muscle contraction) and DNA.
Potassium (K) 0.35 Important component of intracellular fluid and needed to generate the action potential required for muscle contraction.
Sulphur (S) 0.25 Component of vitamins and many proteins.
Sodium (Na) 0.2 Component of extracellular fluid; essential for maintaining water balance and also needed to generate the action potential required for muscle contraction.
Chlorine (Cl) 0.2 Component of extracellular fluid and essential for maintaining water balance.
Magnesium (Mg) 0.1 Important component of many specialized proteins known as enzymes.
Iron (Fe) 0.005 Important component of red blood cells and many enzymes.
small and are impossible for us to see with the naked eye. In fact, approximately 200,000 atoms could fit on the end of a pencil! Atoms are made from the special arrangement of sub- atomic particles known as protons, neutrons and electrons. The basic structure of an atom consists of a central core known as the nucleus in which exist positively charged protons (p+) and uncharged neutrons (n0). Negatively charged electrons (e−) orbit regions around the nucleus known as electron shells, the number of which depends on the specific element.
This basic structure of an atom is shown in Figure 3.1. For ease of understanding, electron shells are depicted diagrammatically as ‘circles’
around the nucleus, although it should be noted that, in reality, electrons do not follow a fixed path or spherical orbits but form negatively charged clouds which surround the nucleus in ‘shells’.
Each electron shell can only hold a specific number of electrons. For example, the first shell (closest to the nucleus) can only ever hold two electrons, while the second can hold eight, the third can hold a maximum of 18 and so on. It
2nd electron shell 1st electron shell nucleus electron proton neutron
Figure 3.1 Basic diagrammatic representation of an atom. In this example, the atom’s nucleus contains six protons and six neutrons, and there are also six electrons in the orbit. As always, the first electron shell holds two electrons and in this case the second shell holds four is important to note that the number of electrons and protons in an atom of an element are always equal. For this reason, the overall charge of an atom is zero.
3.1.3 Atomic number and mass number What makes the atoms of one element different from another is the number of protons present in the nucleus. This is called the atomic number. For example, atoms of hydrogen, oxygen, carbon, nitrogen and so on are primarily different because they each contain different numbers of protons.
The mass number of an atom is the sum of its protons and neutrons. Atoms of elements will always have the same atomic number (i.e. number of protons), but in some cases they may have different numbers of neutrons and hence different mass numbers (see Figure 3.2 for examples). Such atoms are referred to asisotopes. In most cases, these are stable isotopes because their nuclear structure does not change over time. Atoms of the common elements of carbon, nitrogen and oxygen, for example, can exist as stable isotopes. Although isotopes of an element have different numbers of neutrons, they each have identical chemical prop- erties because they all contain the same number of electrons. Progressing from what you learned from Figure 3.1, Figure 3.2 shows how the atomic structure of specific atoms differ from one another on the basis of the number of protons they have.
3.1.4 Atomic mass
The atomic mass (often referred to as atomic weight) of an element is the average mass of all its naturally occurring isotopes. The standard unit for atomic mass (atomic mass unit, abbreviated as ‘amu’) is the dalton, which is symbolized as Da. The mass of a neutron is 1.008 Da, the mass of a proton is 1.007 Da and the mass of an electron is 0.0005 Da. In its purest sense, the Da can be defined as one-twelfth of the mass of a carbon 12C atom and therefore equates to the extremely small mass of 1.66×10−24 grams!
When rounded to the nearest whole number, the atomic mass of an element typically coincides with the mass number of the predominant isotope of that element (see Figure 3.2).
3.1.5 Ions, molecules, compounds and macronutrients
As stated previously, the electrical charge of an atom is neutral because the number of positively charged protons is equal to the number of negatively charged electrons. However, atoms have a characteristic way of becoming charged by gaining or losing one or more electrons. When an atom undergoes this process (calledionization), it now becomes known as anion. More specifically, it will be either an anion (a negatively charged ion because it has gained an electron) or acation (a positively charged ion because it has lost an electron). We can symbolize the ion by writing the chemical symbol of the atom in question followed by the number of positive or negative charges in superscript. For example, Ca2+ designates a calcium ion with two positive charges because it has lost two electrons. Similarly, Cl− designates a chlorine atom (now know as a chloride ion) with a negative charge because it has gained one electron.
Within our bodies, atoms not only exist in free form by themselves but can also join together with other atoms of the same element, or atoms of other elements, to form molecules. A molecule exists when two or more atoms join together.
For example, when two oxygen atoms join
First electron shell
Hydrogen (H) Atomic number = 1 Mass number = 1 or 2 Atomic mass = 1.01
Sodium (Na) Atomic number = 11 Mass number = 23 Atomic mass = 22.99
Atomic number = number of protons in an atom
Mass number = number of protons and neutrons in an atom (boldface indicates most common isotope) Atomic mass = average mass of all stable atoms of a given elements in daltons
Chlorine (Cl) Atomic number = 17 Mass number = 35 or 37 Atomic mass = 35.45
Potassium (K) Atomic number = 19 Mass number = 39, 40, or 41 Atomic mass = 39.10
Iodine (I) Atomic number = 53 Mass number = 127 Atomic mass = 126.90 Carbon (C)
Atomic number = 6 Mass number = 12 or 13 Atomic mass = 12.01
Nitrogen (N) Atomic number = 7 Mass number = 14 or 15 Atomic mass = 14.01
Oxygen (O) Atomic number = 8 Mass number = 16, 17, or 18 Atomic mass = 16.00
Third electron shell
Fourth electron shell
Fifth electron shell Second
electron shell
1p+ 6p+
6n0
7p+ 7n0
8p+ 8n0
53p+ 74n0 19p+
20n0 17p+
18n0 11p+
12n0
Figure 3.2 Basic diagrammatic representation of the atomic structure of several stable atoms. Note the difference in atomic structures between atoms. You should also note that because these atoms can exist as stable isotopes, the mass number can differ because they can differ in the number of neutrons present in the nucleus. (adapted from Tortora and Derrickson,Principles of Anatomy and Physiology, Twelfth Edition, 2009, reproduced by permission of John Wiley & Sons Inc.)
together they form an oxygen molecule which is symbolized as O2.
A compound is a substance which has molecules comprised of atoms of two or more different elements. For example, water is a compound because it consists of two hydrogen atoms joined together with an oxygen atom, and it is thus symbolized as H2O. It is important to note that while all compounds are molecules, not all molecules are compounds.
Most of the chemical elements in our bodies exist in the form of compounds which can be further classified asinorganic compounds (lack- ing carbon) or organic compounds (containing carbon). Some of the most important organic compounds that are relevant tosport and exercise metabolism include carbohydrates, fats and pro- teins. Collectively, these compounds are called the macronutrients, and it is through the action of specific biochemical processes (i.e. chemical
reactions) that our bodies utilize these food sources to provide our muscles with the energy to exercise. Figure 3.3 shows the general flow of chemical organization from atom to macronutrient.