21. 22:
. OU Human Physiology: Chemistry Refresher Introduction . OU Human Physiology: Elements and Atoms: The Building
Blocks of Matter
. OU Human Physiology: Chemical Bonds . OU Human Physiology: Chemical Reactions . OU Human Physiology: Inorganic Compounds Essential to
Human Functioning
. OU Human Physiology: Organic Compounds Essential to
Human Functioning
. OU Human Physiology: Homeostasis Introduction . OU Human Physiology: Overview of Anatomy and Physiology . OU Human Physiology: Structural Organization of the Human
Body
. OU Human Physiology: Functions of Human Life
. OU Human Physiology: Requirements for Human Life . OU Human Physiology: Homeostasis
. OU Human Physiology: Tissue Level Organization
Introduction
. OU Human Physiology: Types of Tissues
. OU Human Physiology: Epithelial Tissue
Protects
. OU Human Physiology: Muscle Tissue and Motion . OU Human Physiology: Nervous Tissue Mediates Perception
and Response
. OU Human Physiology: Introduction to Biomolecules . OU Human Physiology: Organic Compounds Essential to
Human Functioning
OU Human Physiology: Cellular Introduction
OU Human Physiology: The Cytoplasm and Cellular Organelles
23. 24. 25. 26. 2], 28.
29. 30. 31. D2. 33. 34. BD: 36.
a7,
38. 39. AO. 41. 42. 43. 44. 45. A6.
47. 48.
49.
OU Human Physiology: The Nucleus and DNA Replication OU Human Physiology: Protein Synthesis
OU Human Physiology: Cell Growth and Division
OU Human Physiology: Cellular Differentiation
OU Human Physiology: The Cell Membrane Introduction
OU Human Physiology: Structure and Composition of the Cell Membrane
OU Human Physiology: Introduction to Membrane ‘Transport OU Human Physiology: Digestion Introduction
OU Human Physiology: Digestive System Processes OU Human Physiology: The Mouth, Pharynx, and Esophagus OU Human Physiology: The Stomach
OU Human Physiology: The Small and Large Intestines Liver, Pancreas, and Gallbladder
OU Human Physiology: Chemical Digestion and Absorption: A Closer Look
OU Human Physiology: Metabolism Introduction
OU Human Physiology: Overview of Metabolic Reactions OU Human Physiology: Carbohydrate Metabolism
OU Human Physiology: Lipid Metabolism
OU Human Physiology: Protein Metabolism
OU Human Physiology: Metabolic States of the Body
OU Human Physiology: Endocrine Introduction
OU Human Physiology: An Overview of the Endocrine System
OU Human Physiology: Hormones
OU Human Physiology: The Pituitary Gland and Hypothalamus
OU Human Physiology: The Thyroid Gland
50. ol. D2. 33. 04.
Da
36. Of:
58. 59. 60. 61.
62. 63. 64. 65. 66.
67. 68. 69. 70. 71.
72,
73. 74.
OU Human Physiology OU Human Physiology OU Human Physiology OU Human Physiology
Endocrine System OU Human Physiology OU Human Physiology Nervous System
OU Human Physiology OU Human Physiology OU Human Physiology
OU Human Physiology:
Integration OU Human Physiology OU Human Physiology OU Human Physiology
: The Parathyroid Glands : The Adrenal Glands
: The Pineal Gland
: The Endocrine Pancreas
: Neural Communication Introduction : Basic Structure and Function of the
: Nervous Tissue : The Function of Nervous Tissue : The Action Potential
Synaptic Transmission and Neural
: The Nervous System Introduction : The Embryologic Perspective : The Central Nervous System
OU Human Physiology: Special Senses and Reflexes
Introduction
: Central Processing
OU Human Physiology OU Human Physiology OU Human Physiology OU Human Physiology Nervous System
OU Human Physiology System
OU Human Physiology OU Human Physiology
: Motor Responses : Reflexes : Introduction to the Autonomic
: Divisions of the Autonomic Nervous
: Autonomic Reflexes and Homeostasis - Central Control
7D:
76.
Ta.
78. 79:
80.
81. 82. 83. 84. 85.
86. 87.
88. 89. 90. 91. o2,
93.
94. 9D:
96. 97. 98.
OU Human Physiology: Drugs that Affect the Autonomic
system OU Human Physiology
OU Human Physiology OU Human Physiology Relaxation
OU Human Physiology Tension
OU Human Physiology:
: Muscle Tissue Introduction . OU Human Physiology: Overview of Muscle Tissues : Skeletal Muscle
- Muscle Fiber Contraction and
: Nervous System Control of Muscle
Types of Muscle Fibers
OU Human Physiology OU Human Physiology OU Human Physiology OU Human Physiology Muscle Tissue
OU Human Physiology:
: Exercise and Muscle Performance : Smooth Muscle
: Cardiac Muscle Tissue
: Development and Regeneration of
Respiratory System Introduction
OU Human Physiology: Organs and Structures of the
Respiratory System OU Human Physiology: OU Human Physiology: OU Human Physiology:
The Lungs The Process of Breathing Gas Exchange
OU Human Physiology: Transport of Gases
Functions OU Human Physiology: Introduction
OU Human Physiology: OU Human Physiology: Activity
OU Human Physiology: OU Human Physiology: OU Human Physiology:
The Cardiovascular System Heart:
Heart Anatomy Cardiac Muscle and Electrical
Cardiac Cycle Cardiac Physiology Blood Vessels and Blood Introduction
99.
100.
101. 102.
103. 104. 105. 106. 107. 108. 109. 110. 111. 112. 113. 114. 115. 116. 117.
118. 119.
OU Human Physiology: Vessels
Resistance
Structure and Function of Blood
OU Human Physiology: Capillary Exchange
OU Human Physiology Vascular System
OU Human Physiology OU Human Physiology OU Human Physiology OU Human Physiology:
OU Human Physiology OU Human Physiology OU Human Physiology
Immune Systems OU Human Physiology: Immune Response
lymphocytes and Antibodies The Immune Response against
OU Human Physiology: Pathogens OU Human Physiology:
: Homeostatic Regulation of the
: Blood Introduction : An Overview of Blood : Production of the Formed Elements
Erythrocytes
: Leukocytes and Platelets : Hemostasis : Blood Typing
Diseases Associated with Depressed
or Overactive Immune Responses OU Human Physiology: Transplantation and Cancer
Immunology OU Human Physiology:
Introduction to the Urinary System
OU Human Physiology: Physical Characteristics of Urine
120. 121. 122, 123. 124. 125. 126.
127,
128. 129.
130. 131. 132. 133. 134. 135. 136.
137.
OU Human Physiology: Physiology of Urine Formation
OU Human Physiology: Tubular Reabsorption
OU Human Physiology: Regulation of Renal Blood Flow Function
OU Human Physiology: Regulation of Fluid Volume and Composition
OU Human Physiology: The Urinary System and Homeostasis
Introduction
OU Human Physiology: Anatomy and Physiology of the Male Reproductive System
OU Human Physiology: Anatomy and Physiology of the Female Reproductive System
OU Human Physiology: Development of the Male and Female Reproductive Systems
OU Human Physiology: Development Introduction
OU Human Physiology: Fertilization
Labor, and Birth OU Human Physiology: Lactation
OU Human Physiology: Chemistry Refresher Introduction class="introduction" Human DNA
Human DNA is described as a double
helix that resembles a molecular spiral Staircase. In humans the DNA is organized into 46 chromosomes
Note: Chapter Objectives After studying this chapter, you will be able to:
Describe the fundamental composition of matter
Identify the three subatomic particles
Identify the four most abundant elements in the body
Explain the relationship between an atom’s number of electrons and its relative stability
Distinguish between ionic bonds, covalent bonds, and hydrogen bonds Explain how energy is invested, stored, and released via chemical reactions, particularly those reactions that are critical to life
Explain the importance of the inorganic compounds that contribute to life, such as water, salts, acids, and bases
e¢ Compare and contrast the four important classes of organic (carbon- based) compounds—proteins, carbohydrates, lipids and nucleic acids —according to their composition and functional importance to human life
The smallest, most fundamental material components of the human body are basic chemical elements. In fact, chemicals called nucleotide bases are the foundation of the genetic code with the instructions on how to build and maintain the human body from conception through old age. There are about three billion of these base pairs in human DNA.
Human chemistry includes organic molecules (carbon-based) and biochemicals (those produced by the body). Human chemistry also includes elements. In fact, life cannot exist without many of the elements that are part of the earth. All of the elements that contribute to chemical reactions, to the transformation of energy, and to electrical activity and muscle contraction—elements that include phosphorus, carbon, sodium, and calcium, to name a few—originated in stars.
These elements, in turn, can form both the inorganic and organic chemical compounds important to life, including, for example, water, glucose, and proteins. This chapter begins by examining elements and how the structures of atoms, the basic units of matter, determine the characteristics of elements by the number of protons, neutrons, and electrons in the atoms. The chapter then builds the framework of life from there.
OU Human Physiology: Elements and Atoms: The Building Blocks of Matter By the end of this section, you will be able to:
e Discuss the relationships between matter, mass, elements, compounds, atoms, and subatomic particles
e Distinguish between atomic number and mass number
e Identify the key distinction between isotopes of the same element
e Explain how electrons occupy electron shells and their contribution to an atom’s relative stability
The substance of the universe—from a grain of sand to a star—is called matter. Scientists define matter as anything that occupies space and has mass. An object’s mass and its weight are related concepts, but not quite the same. An object’s mass is the amount of matter contained in the object, and the object’s mass is the same whether that object is on Earth or in the zero- gravity environment of outer space. An object’s weight, on the other hand, is its mass as affected by the pull of gravity. Where gravity strongly pulls on an object’s mass its weight is greater than it is where gravity is less strong. An object of a certain mass weighs less on the moon, for example, than it does on Earth because the gravity of the moon is less than that of Earth. In other words, weight is variable, and is influenced by gravity. A piece of cheese that weighs a pound on Earth weighs only a few ounces on the moon.
Elements and Compounds
All matter in the natural world is composed of one or more of the 92 fundamental substances called elements. An element is a pure substance that is distinguished from all other matter by the fact that it cannot be created or broken down by ordinary chemical means. While your body can assemble many of the chemical compounds needed for life from their constituent elements, it cannot make elements. They must come from the environment. A familiar example of an element that you must take in is calcium (Ca**). Calcium is essential to the human body; it is absorbed and used for a number of processes, including strengthening bones. When you consume dairy products your digestive system breaks down the food into
components small enough to cross into the bloodstream. Among these is calcium, which, because it is an element, cannot be broken down further. The elemental calcium in cheese, therefore, is the same as the calcium that forms your bones. Some other elements you might be familiar with are oxygen, sodium, and iron. The elements in the human body are shown in [link], beginning with the most abundant: oxygen (O), carbon (C), hydrogen (H), and nitrogen (N). Each element’s name can be replaced by a one- or two-letter symbol; you will become familiar with some of these during this course. All the elements in your body are derived from the foods you eat and the air you breathe.
Elements of the Human Body
Oxygen
cobalt (Co), copper (Cu), fluorine ‘P), iodine (I), iron (Fe), manganese (Mn), molybdenum (Mo), selenium (Se), silicon (Si), tin (Sn), vanadium and zinc (Zn).
Hydrogen
Carbon
The main elements that compose the human body are shown from most abundant to least abundant.
In nature, elements rarely occur alone. Instead, they combine to form compounds. A compound is a substance composed of two or more elements joined by chemical bonds. For example, the compound glucose is an important body fuel. It is always composed of the same three elements: carbon, hydrogen, and oxygen. Moreover, the elements that make up any given compound always occur in the same relative amounts. In glucose,
there are always six carbon and six oxygen units for every twelve hydrogen units. But what, exactly, are these “units” of elements?
Atoms and Subatomic Particles
An atom is the smallest quantity of an element that retains the unique properties of that element. In other words, an atom of hydrogen is a unit of hydrogen—the smallest amount of hydrogen that can exist. As you might guess, atoms are almost unfathomably small. The period at the end of this sentence is millions of atoms wide.
Atomic Structure and Energy
Atoms are made up of even smaller subatomic particles, three types of which are important: the proton, neutron, and electron. The number of positively-charged protons and non-charged (“neutral”) neutrons, gives mass to the atom, and the number of each in the nucleus of the atom determine the element. The number of negatively-charged electrons that “spin” around the nucleus at close to the speed of light equals the number of protons. An electron has about 1/2000th the mass of a proton or neutron.
[link] shows two models that can help you imagine the structure of an atom —in this case, helium (He). In the planetary model, helium’s two electrons are shown circling the nucleus in a fixed orbit depicted as a ring. Although this model is helpful in visualizing atomic structure, in reality, electrons do not travel in fixed orbits, but whiz around the nucleus erratically in a so- called electron cloud.
Two Models of Atomic Structure
Electron
(a) Planetary model
Nucleus
Cloud of negative charge (2 electrons)
(b) Electron cloud model
(a) In the planetary model, the electrons of helium are shown in fixed orbits, depicted as rings, at a precise distance from the nucleus, somewhat like planets orbiting the sun. (b) In the electron cloud model, the electrons of carbon are shown in the variety of locations they would have at different distances from the nucleus over time.
An atom’s protons and electrons carry electrical charges. Protons, with their positive charge, are designated p*. Electrons, which have a negative charge,
are designated e-. An atom’s neutrons have no charge: they are electrically neutral. Just as a magnet sticks to a steel refrigerator because their opposite charges attract, the positively charged protons attract the negatively charged electrons. This mutual attraction gives the atom some structural stability. The attraction by the positively charged nucleus helps keep electrons from straying far. The number of protons and electrons within a neutral atom are equal, thus, the atom’s overall charge is balanced.
Atomic Number and Mass Number
An atom of carbon is unique to carbon, but a proton of carbon is not. One proton is the same as another, whether it is found in an atom of carbon, sodium (Na), or iron (Fe). The same is true for neutrons and electrons. So, what gives an element its distinctive properties—what makes carbon so different from sodium or iron? The answer is the unique quantity of protons each contains. Carbon by definition is an element whose atoms contain six protons. No other element has exactly six protons in its atoms. Moreover, all atoms of carbon, whether found in your liver or in a lump of coal, contain six protons. Thus, the atomic number, which is the number of protons in the nucleus of the atom, identifies the element. Because an atom usually has the same number of electrons as protons, the atomic number identifies the usual number of electrons as well.
In their most common form, many elements also contain the same number of neutrons as protons. The most common form of carbon, for example, has six neutrons as well as six protons, for a total of 12 subatomic particles in its nucleus. An element’s mass number is the sum of the number of protons and neutrons in its nucleus. So the most common form of carbon’s mass number is 12. (Electrons have so little mass that they do not appreciably contribute to the mass of an atom.) Carbon is a relatively light element. Uranium (U), in contrast, has a mass number of 238 and is referred to as a heavy metal. Its atomic number is 92 (it has 92 protons) but it contains 146 neutrons; it has the most mass of all the naturally occurring elements.
The periodic table of the elements, shown in [link], is a chart identifying the 92 elements found in nature, as well as several larger, unstable elements discovered experimentally. The elements are arranged in order of their atomic number, with hydrogen and helium at the top of the table, and the more massive elements below. The periodic table is a useful device because for each element, it identifies the chemical symbol, the atomic number, and the mass number, while organizing elements according to their propensity to react with other elements. The number of protons and electrons in an element are equal. The number of protons and neutrons may be equal for some elements, but are not equal for all.
The Periodic Table of the Elements
PERIODIC TABLE
NIsT
Group . National Institute of . the Element se 1 Atomic Properties of the Elements standard ond Tehnoloay 18 IA VIIA Frequently used fundamental physical constants Physics Standard ly Pl For the most accurate values of these and other constants, vsit physics.rist goviconstants, Laboratory Reference Data 1 4 second = 9 192 631 770 periods of radiation corresponding to the ansition physics.nist.gov www.nist.gov/std etween the two hyperfine levels e ground state of S 2 speed of light in vacuum 299792458 ms’ —_ (exact) LJ Solids 13 14 15 16 17 A Planck constant h 6.6261 x toMy s (n= hi2n) GD Liquids IIA IVA VA VIA VIIA 3 *s 14 Is, elementary charge 6 1.6022 x 10°C 5 “p16 Sp Li Me B. ) electron mass m, 9.1094 x 10°" kg Gases 2 Cc ® 2 in ‘4 e mac? 0.5110 MeV, C Artificially ithium jeryllium proton mass m, 1.6726 x 10°” kg Prepared Boron Carbon Brie Oe 2 fine-structure constant a 11137.036 ye oo ‘ Bee <"2s 1s'2s a s'25°2n | 18%25%29° s3917_| 9.3297 Rekayeemars pee bred T aies 2980 | _ 11.2603 % 7 = . 43 p>] Es 4ge 3 11 *s,,/12 ‘s, Rihc 13.6057 eV 13 *Pi2|14 P,]15 _‘s;,|16 °P, Na Boltzmann constant k 4.3807 x 107°) «7 Al l Ss 3 ium | Magnesium i Aluminum | Silicon | Phosphorus | — Suffur 22.98976928 4.3050 3 4 5 6 F i 8 Q 4 0 4 1 12 26.9815386 28.0855 30. 973762 32.065 (Ne}3s INel3s" 3 b mM INe}as"3p | [Ne}3s°3p? | [Nej3sap* | INeT39°ap* 1 l 5.1391 7.8462 WB vB vB VIB VIB w= IB IIB 5.9868 groi7__| 10.4867 | 10.3600 19 ’s,,/20 ‘'s,/21 *D,,/22 *F,|23 ‘F,,|24 's,|25 °s,.]26 °D,)27 ‘F,.}28 F,|29 ’s,,/30 'S./31 *P?.|/32_ *P,|33 ‘si,|34 *, 3 Ca | Sc | Ti Cr | Mn | Fe | Co | Ni u | Zn | Ga | Ge | As | Se © 4) potassium | Calcium | Scandium | Tlanium | Vanadium | Chromium | Manganese | Iron Cobalt Nickel Copper Zine Gallium | Germanium | Arsenic | Selenium F per 39.0883 | 40.078 | 44955912] 47.867 | 50.9415 | 51.9961 | 54938045 | 55.845 | 58.933195 | 58.6934 | 63.546 65.38, 69.723 72.64 | 7492160 | 78.98 o 2 2 242 2 4, B42 10, 10402 104 2, ‘ [Ariés [Aras’ [araaas’ [Ar3d?4s: [arad"4s farjad’4s | jarad’4s” | [arizd’4s° | [Araa’4s” | [anad®as’ tansa'4s | parad'ae? | farjsd™4s"4p | pariad’“4s%4p” | (arjad™”4s"4p* | farjad™4s"40* 4.3407 6.1132 6.5615 6.8284 6.7462 6.7665 7.434 7.9024 7.8810 7.6399 7.7264 9.3942 5.9993 7.8994 9.7886 9.7524 37 _ *S,, 5 ‘Sy ay” a =, “Nb” “M 'S, “R *F5/45 *F,.}46 's,/47 7*s,,/48 's,|/49 *Pi,|50 *P,/51_ ‘si,|52_ *P,|53 _*P5, r r ° u Cc n| Sn | Sb | Te 5] Rubidium | strontiom | yttium | Zirconium | — Niobium | Molyodenum Ruthenium | Rhodium | Palladium | — Siver | Cadmium | Indium Tin Antimony | Tellurium lodine 85.4678 87.62 3.90586 | 91.224 | 92.9638 95.96 1o2.90650 | 106.42 | to7ses2 | 11241 | 114818 | 18.710 | 121.760 127.60 | 126.90447_ IKr}5s Ikri5s° Ukri4a5s? | [Kri4a’ss’ [kr}4d*ss [krl4a°5s [kry4e’5s [kr]4d°5s [Krad ikrj4d'5s | [kri4d “5s? | [kr]4d"5s"5p | kr}4d“5s”5p" | 1KrJ4e™5s%5p" | Kr]4d°5s"5p* I tkri4d'75s%5p° 4771 5.6949 6.2173 | 6.6339 6.7589 7.0924 7.3605 7.4589 8.3369 7.5762 8.9938 5.7864 7.3439 8.5084 9.0096 | 10.4513, 55 °*s,,/56 ‘s, 72 76,)73 “Feo| 74 °D,|75 “Seo} 76 *D,|77_ “Fy. 78 °D,/79 *S,2] 81 *Pi,/82 *P,/83 ‘s;,|84 *P,|85 “Ps, Cs a Hf | Ta Re | Os | Ir | Pt | Au Tl | Pb | Bi | Po| At 6} cesium Barium Hafnium | Tantalum | Tungsten | Rhenium | Osmium Iridium | Platinum Thallium Lead Bismuth | Polonium | Astatine 192.9054519 | 137.327 178.49 | 180.94788 | 183.84 186.207 | 190.23 | 192.217 | 195.084 | 198.966569 204.3833 | 207.2 | 20898040 | (208) (210) [Xe]és [xe]6s” [xe]4t''5d’6s”| (xe}4t"*5a°6s” | [xe]4*"*50“6e | (Xe]4t"“5d°6s” | [Xe]4t"*5d°6s" | [xe]4t''5a”6s"| [xelat!*5d°6s | [xe}41"*5d""6s, [Hg]sp [Hal6p” IHal6p° [Hal6p* IHgi6p® 3.8939 5.2117 6.8251 7.5496 7.8640 7. 8. 2 8.9670 8.9588 9.2255 6.1082 7.4167 7.2855 8.414 87 ’s,,/88 's, Fr a 7] Francium | Radium (223) (226) [Roj7s [Rnj7s? 4.0727 5.2784 Atomic Ground-state 2/57 *D..|58 ‘G:|59 ‘T3,|60 69 *F, Number Level q FS ! 2| La e | Pr | Nd Tm ‘Symbol | Lanthanum | Cerium — [Praseodymium) Neodymium | P Samarium | Europlum Terbium | Dysprosium Lutetium | 19890547 | 140.116 | +40.90765 | 144.242 150.36 | 151.964 158.92535 | 162.500 168.93421 174.9668 | [xejsdes? | [xeystsaes? | [xelar’es? | xel4ttes” [xeyat’es? | [xeyss es? prelates” | (xelat"6s* | [xejer"'6s” pxeyar'%6s* | pxoyar'Mes? | preyer!*sdes” Name 5.5769 5.5387 5.475 5.5250 5.6437 5.6704 5.8638 5.9389 | 6.0215 6.1843 6.2542 5.4258 89 *0,,/90 *F,|91 *,,,|92 Ly a; 8 ; Atomic. g ‘ Weight" 3| Ac | Th a | Actinium | Thorium | Protactnium | Uranium @ nf jonizatio qt (221): Z82,05806 ere aetaan? aginaatis ee [Rnj6d7s° | [Rnj6d’7s° | [Rnjsf°6d7s° | [RnI6f'ed7s Configuration Energy (eV) 5.3807 6.3067 5.89 6.1939
‘Based upon "°C. () indicates the mass number of the longest-lived isotope. »f the data, visit physics.nist.gov/data NIST SP 966 (September 2010)
(credit: R.A. Dragoset, A. Musgrove, C.W. Clark, W.C. Martin)
Note:
Visit this website to view the periodic table. In the periodic table of the elements, elements in a single column have the same number of electrons that can participate in a chemical reaction. These electrons are known as “valence electrons.” For example, the elements in the first column all have a single valence electron, an electron that can be “donated” in a chemical reaction with another atom. What is the meaning of a mass number shown in parentheses?
Isotopes
Although each element has a unique number of protons, it can exist as different isotopes. An isotope is one of the different forms of an element, distinguished from one another by different numbers of neutrons. The standard isotope of carbon is '*C, commonly called carbon twelve. !*C has six protons and six neutrons, for a mass number of twelve. All of the isotopes of carbon have the same number of protons; therefore, °C has seven neutrons, and !4C has eight neutrons. The different isotopes of an element can also be indicated with the mass number hyphenated (for example, C-12 instead of '*C). Hydrogen has three common isotopes, shown in [link].
Isotopes of Hydrogen
eo )(@ B Q / eQ | Q
Protium (1H) Deuterium (7H) Tritium (3H)
Protium, designated 1H, has one proton and no neutrons. It is by far the most abundant isotope of hydrogen in nature. Deuterium, designated 2H, has one proton and one neutron. Tritium, designated 5H, has two neutrons.
An isotope that contains more than the usual number of neutrons is referred to as a heavy isotope. An example is '4C. Heavy isotopes tend to be unstable, and unstable isotopes are radioactive. A radioactive isotope is an isotope whose nucleus readily decays, giving off subatomic particles and electromagnetic energy. Different radioactive isotopes (also called radioisotopes) differ in their half-life, the time it takes for half of any size sample of an isotope to decay. For example, the half-life of tritium—a radioisotope of hydrogen—is about 12 years, indicating it takes 12 years for half of the tritium nuclei in a sample to decay. Excessive exposure to radioactive isotopes can damage human cells and even cause cancer and birth defects, but when exposure is controlled, some radioactive isotopes can be useful in medicine. For more information, see the Career Connections.
Note:
Career Connection
Interventional Radiologist
The controlled use of radioisotopes has advanced medical diagnosis and treatment of disease. Interventional radiologists are physicians who treat disease by using minimally invasive techniques involving radiation. Many conditions that could once only be treated with a lengthy and traumatic operation can now be treated non-surgically, reducing the cost, pain, length of hospital stay, and recovery time for patients. For example, in the past, the only options for a patient with one or more tumors in the liver were surgery and chemotherapy (the administration of drugs to treat cancer). Some liver tumors, however, are difficult to access surgically, and others
could require the surgeon to remove too much of the liver. Moreover, chemotherapy is highly toxic to the liver, and certain tumors do not respond well to it anyway. In some such cases, an interventional radiologist can treat the tumors by disrupting their blood supply, which they need if they are to continue to grow. In this procedure, called radioembolization, the radiologist accesses the liver with a fine needle, threaded through one of the patient’s blood vessels. The radiologist then inserts tiny radioactive “seeds” into the blood vessels that supply the tumors. In the days and weeks following the procedure, the radiation emitted from the seeds destroys the vessels and directly kills the tumor cells in the vicinity of the treatment.
Radioisotopes emit subatomic particles that can be detected and tracked by imaging technologies. One of the most advanced uses of radioisotopes in medicine is the positron emission tomography (PET) scanner, which detects the activity in the body of a very small injection of radioactive glucose, the simple sugar that cells use for energy. The PET camera reveals to the medical team which of the patient’s tissues are taking up the most glucose. Thus, the most metabolically active tissues show up as bright “hot spots” on the images ([link]). PET can reveal some cancerous masses because cancer cells consume glucose at a high rate to fuel their rapid reproduction.
PET Scan
PET highlights areas in the body where there is relatively high glucose use, which is characteristic of cancerous tissue. This PET scan shows sites of the spread of a large primary tumor to other sites.
The Behavior of Electrons
In the human body, atoms do not exist as independent entities. Rather, they are constantly reacting with other atoms to form and to break down more complex substances. To fully understand anatomy and physiology you must grasp how atoms participate in such reactions. The key is understanding the behavior of electrons.
Although electrons do not follow rigid orbits a set distance away from the atom’s nucleus, they do tend to stay within certain regions of space called
electron shells. An electron shell is a layer of electrons that encircle the nucleus at a distinct energy level.
The atoms of the elements found in the human body have from one to five electron shells, and all electron shells hold eight electrons except the first shell, which can only hold two. This configuration of electron shells is the same for all atoms. The precise number of shells depends on the number of electrons in the atom. Hydrogen and helium have just one and two electrons, respectively. If you take a look at the periodic table of the elements, you will notice that hydrogen and helium are placed alone on either sides of the top row; they are the only elements that have just one electron shell ({link]). A second shell is necessary to hold the electrons in all elements larger than hydrogen and helium.
Lithium (Li), whose atomic number is 3, has three electrons. Two of these fill the first electron shell, and the third spills over into a second shell. The second electron shell can accommodate as many as eight electrons. Carbon, with its six electrons, entirely fills its first shell, and half-fills its second. With ten electrons, neon (Ne) entirely fills its two electron shells. Again, a look at the periodic table reveals that all of the elements in the second row, from lithium to neon, have just two electron shells. Atoms with more than ten electrons require more than two shells. These elements occupy the third and subsequent rows of the periodic table.
Electron Shells
Electron
(b) (c)
Electrons orbit the atomic nucleus at distinct levels of energy called electron shells. (a) With one electron, hydrogen only half-fills its electron shell. Helium also has a single shell, but its two electrons completely fill it. (b) The electrons of carbon completely fill its first electron shell, but only half-fills its second. (c) Neon, an element that does not occur in the body, has 10 electrons, filling both of its electron shells.
The factor that most strongly governs the tendency of an atom to participate in chemical reactions is the number of electrons in its valence shell. A valence shell is an atom’s outermost electron shell. If the valence shell is full, the atom is stable; meaning its electrons are unlikely to be pulled away from the nucleus by the electrical charge of other atoms. If the valence shell is not full, the atom is reactive; meaning it will tend to react with other
atoms in ways that make the valence shell full. Consider hydrogen, with its one electron only half-filling its valence shell. This single electron is likely to be drawn into relationships with the atoms of other elements, so that hydrogen’s single valence shell can be stabilized.
All atoms (except hydrogen and helium with their single electron shells) are most stable when there are exactly eight electrons in their valence shell. This principle is referred to as the octet rule, and it states that an atom will give up, gain, or share electrons with another atom so that it ends up with eight electrons in its own valence shell. For example, oxygen, with six electrons in its valence shell, is likely to react with other atoms in a way that results in the addition of two electrons to oxygen’s valence shell, bringing the number to eight. When two hydrogen atoms each share their single electron with oxygen, covalent bonds are formed, resulting in a molecule of water, H»O.
In nature, atoms of one element tend to join with atoms of other elements in characteristic ways. For example, carbon commonly fills its valence shell by linking up with four atoms of hydrogen. In so doing, the two elements form the simplest of organic molecules, methane, which also is one of the most abundant and stable carbon-containing compounds on Earth. As stated above, another example is water; oxygen needs two electrons to fill its valence shell. It commonly interacts with two atoms of hydrogen, forming HO. Incidentally, the name “hydrogen” reflects its contribution to water (hydro- = “water”; -gen = “maker”). Thus, hydrogen is the “water maker.”
Chapter Review
The human body is composed of elements, the most abundant of which are oxygen (O), carbon (C), hydrogen (H) and nitrogen (N). You obtain these elements from the foods you eat and the air you breathe. The smallest unit of an element that retains all of the properties of that element is an atom. But, atoms themselves contain many subatomic particles, the three most important of which are protons, neutrons, and electrons. These particles do not vary in quality from one element to another; rather, what gives an element its distinctive identification is the quantity of its protons, called its atomic number. Protons and neutrons contribute nearly all of an atom’s
mass; the number of protons and neutrons is an element’s mass number. Heavier and lighter versions of the same element can occur in nature because these versions have different numbers of neutrons. Different versions of an element are called isotopes.
The tendency of an atom to be stable or to react readily with other atoms is largely due to the behavior of the electrons within the atom’s outermost electron shell, called its valence shell. Helium, as well as larger atoms with eight electrons in their valence shell, is unlikely to participate in chemical reactions because they are stable. All other atoms tend to accept, donate, or share electrons in a process that brings the electrons in their valence shell to eight (or in the case of hydrogen, to two).
Glossary
atom smallest unit of an element that retains the unique properties of that element
atomic number number of protons in the nucleus of an atom
compound substance composed of two or more different elements joined by chemical bonds
electron subatomic particle having a negative charge and nearly no mass; found orbiting the atom’s nucleus
electron shell area of space a given distance from an atom’s nucleus in which electrons are grouped
element substance that cannot be created or broken down by ordinary chemical means
isotope one of the variations of an element in which the number of neutrons differ from each other
mass number sum of the number of protons and neutrons in the nucleus of an atom
matter physical substance; that which occupies space and has mass
neutron heavy subatomic particle having no electrical charge and found in the atom’s nucleus
periodic table of the elements arrangement of the elements in a table according to their atomic number; elements having similar properties because of their electron arrangements compose columns in the table, while elements having the same number of valence shells compose rows in the table
proton heavy subatomic particle having a positive charge and found in the atom’s nucleus
radioactive isotope unstable, heavy isotope that gives off subatomic particles, or electromagnetic energy, as it decays; also called radioisotopes
valence shell outermost electron shell of an atom
OU Human Physiology: Chemical Bonds By the end of this section, you will be able to:
e Explain the relationship between molecules and compounds e Distinguish between ions, cations, and anions
e Identify the key difference between ionic and covalent bonds e Distinguish between nonpolar and polar covalent bonds
e Explain how water molecules link via hydrogen bonds
Atoms separated by a great distance cannot link; rather, they must come close enough for the electrons in their valence shells to interact. But do atoms ever actually touch one another? Most physicists would say no, because the negatively charged electrons in their valence shells repel one another. No force within the human body—or anywhere in the natural world—is strong enough to overcome this electrical repulsion. So when you read about atoms linking together or colliding, bear in mind that the atoms are not merging in a physical sense.
Instead, atoms link by forming a chemical bond. A bond is a weak or strong electrical attraction that holds atoms in the same vicinity. The new grouping is typically more stable—less likely to react again—than its component atoms were when they were separate. A more or less stable grouping of two or more atoms held together by chemical bonds is called a molecule. The bonded atoms may be of the same element, as in the case of Hp, which is called molecular hydrogen or hydrogen gas. When a molecule is made up of two or more atoms of different elements, it is called a chemical compound. Thus, a unit of water, or H»O, is a compound, as is a single molecule of the gas methane, or CHy.
Three types of chemical bonds are important in human physiology, because they hold together substances that are used by the body for critical aspects of homeostasis, signaling, and energy production, to name just a few important processes. These are ionic bonds, covalent bonds, and hydrogen bonds.
Ions and Ionic Bonds
Recall that an atom typically has the same number of positively charged protons and negatively charged electrons. As long as this situation remains, the atom is electrically neutral. But when an atom participates in a chemical reaction that results in the donation or acceptance of one or more electrons, the atom will then become positively or negatively charged. This happens frequently for most atoms in order to have a full valence shell, as described previously. This can happen either by gaining electrons to fill a shell that is more than half-full, or by giving away electrons to empty a shell than is less than half-full, thereby leaving the next smaller electron shell as the new, full, valence shell. An atom that has an electrical charge—whether positive or negative—is an ion.
Note:
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— ee COLLEGE
Visit this website to learn about electrical energy and the attraction/repulsion of charges. What happens to the charged electroscope when a conductor is moved between its plastic sheets, and why?
Potassium (K), for instance, is an important element in all body cells. Its atomic number is 19. It has just one electron in its valence shell. This characteristic makes potassium highly likely to participate in chemical reactions in which it donates one electron. (It is easier for potassium to donate one electron than to gain seven electrons.) The loss will cause the positive charge of potassium’s protons to be more influential than the negative charge of potassium’s electrons. In other words, the resulting potassium ion will be slightly positive. A potassium ion is written K*, indicating that it has lost a single electron. A positively charged ion is known as a Cation.
Now consider fluorine (F), a component of bones and teeth. Its atomic number is nine, and it has seven electrons in its valence shell. Thus, it is highly likely to bond with other atoms in such a way that fluorine accepts one electron (it is easier for fluorine to gain one electron than to donate seven electrons). When it does, its electrons will outnumber its protons by one, and it will have an overall negative charge. The ionized form of fluorine is called fluoride, and is written as F~. A negatively charged ion is known as an anion.
Atoms that have more than one electron to donate or accept will end up with stronger positive or negative charges. A cation that has donated two electrons has a net charge of +2. Using magnesium (Mg) as an example, this can be written Mg** or Mg?*. An anion that has accepted two electrons has a net charge of —2. The ionic form of selenium (Se), for example, is typically written Se*.
The opposite charges of cations and anions exert a moderately strong mutual attraction that keeps the atoms in close proximity forming an ionic bond. An ionic bond is an ongoing, close association between ions of opposite charge. The table salt you sprinkle on your food owes its existence to ionic bonding. As shown in [link], sodium commonly donates an electron to chlorine, becoming the cation Na*. When chlorine accepts the electron, it becomes the chloride anion, Cl-. With their opposing charges, these two ions strongly attract each other.
Ionic Bonding
Net positive charge
(a) Sodium readily donates the solitary electron in its valence shell to chlorine, which needs only one electron to have a full valence shell. (b) The opposite electrical charges of the resulting sodium cation and chloride anion result in the formation of a bond of attraction called an ionic bond. (c) The attraction of
many sodium and chloride ions results in the formation of large groupings called crystals.
Water is an essential component of life because it is able to break the ionic bonds in salts to free the ions. In fact, in biological fluids, most individual atoms exist as ions. These dissolved ions produce electrical charges within the body. The behavior of these ions produces the tracings of heart and brain function observed as waves on an electrocardiogram (EKG or ECG) or an electroencephalogram (EEG). The electrical activity that derives from the interactions of the charged ions is why they are also called electrolytes.
Covalent Bonds
Unlike ionic bonds formed by the attraction between a cation’s positive charge and an anion’s negative charge, molecules formed by a covalent bond share electrons in a mutually stabilizing relationship. Like next-door neighbors whose kids hang out first at one home and then at the other, the atoms do not lose or gain electrons permanently. Instead, the electrons move back and forth between the elements. Because of the close sharing of pairs of electrons (one electron from each of two atoms), covalent bonds are stronger than ionic bonds.
Nonpolar Covalent Bonds
[link] shows several common types of covalent bonds. Notice that the two covalently bonded atoms typically share just one or two electron pairs, though larger sharings are possible. The important concept to take from this is that in covalent bonds, electrons in the outermost valence shell are shared to fill the valence shells of both atoms, ultimately stabilizing both of the