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Chapter 3: e Chemical Level of Organization
Chapter Introduction
e smallest, most fundamental material components of the human body are basic chemical elements. An element is a pure substance containing only the same type of atom.
You may have studied chemist in a dedicated course, but here we will examine human chemist . is study is somewhat distinct because we will examine organic (carbon-based) molecules, inorganic (non carbon-based) molecules, and biochemicals (those produced by the body) and all of the chemical reactions we examine take place in an aqueous (wate ) environment, rather than in air.
is chapter begins by examining elements and how the structures of the smallest units of matter determine the characteristics of the elements.
3.1Elements and Atoms: e Building Blocks of Matter
e substance that makes up the universe—whether we are talking about a grain of sand, a star in the sky, or a potato chip—is called matter. Scientists de ne matter as anything that occupies space and has mass. An object’s mass and its weight are related concepts, but not quite the same. Mass is the amount of matter contained in an object, and the object’s mass is the same whether that object is on Ea h or in the zero-gravity environment of outer space. An object’s weight, on the other hand, is its mass as a ected by the pull of gravity. For example, a piece of cheese that weighs a pound on Ea h weighs only a few ounces on the moon.
3.1aElements and Compounds
All matter in the natural world is composed of one or more of the 92 fundamental substances called elements. While your body can assemble many of the chemical compounds needed for life from their constituent elements, it cannot make elements. ey must come from the environment. A familiar example of an element that you must take in is calcium. Calcium is essential to the human body; it is absorbed and used for a number of processes, including muscle contraction and strengthening bones. Imagine you are snacking on a slice of cheese or a dish of lentils; your digestive system breaks down the food into components small enough to cross into the bloodstream (these processes are discussed in Chapter 25 ). Among these components, calcium, as an element, cannot be broken down fu her. e calcium in your food, therefore, is the same as the calcium that forms your bones. Some other elements you might be familiar with are oxygen, sodium, and iron. e elements in the human body are shown in Figure 3.1, beginning with the most abundant: oxygen, carbon, hydrogen, and nitrogen. All the elements in your body are derived from the foods you eat and the air you breathe.
misconception more accurately by drawing in the electron cloud. In this chapter we will use the format of depicting the electrons in a xed ring.
Figure 3.2Two Models of Atomic Structure
(A) In this model, the electrons of helium are shown in xed orbits, depicted as rings, at a precise distance from the nucleus, somewhat like planets orbiting the sun. (B) In this more accurate model, the electrons of carbon are shown in the variety of locations they would have at di erent distances from the nucleus over time.
Atomic Number and Mass Number
If atoms are composed of protons, neutrons, and electrons, and one proton is the same as another, whether it is found in an atom of carbon, sodium, or iron, then what gives an element its distinctive prope ies—what makes carbon so di erent from sodium or iron? e answer is the unique quantity of protons each contains. Carbon, by de nition, 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. us, the atomic number, the number of protons in the nucleus of the atom, identi es the element. Because an atom usually has the same number of electrons as protons, the atomic number identi es the usual number of electrons as well. e atomic weight of the atom is its weight, measured in atomic mass units (amu) Because electrons have so much less mass than protons and neutrons, the atomic weight is roughly equal to the number of protons and neutrons plus a little weight from the electrons. ese numbers, as well as the symbol and name of each element can be seen in the periodic table of elements (Figure 3.3A). e elements are arranged in order of their atomic number, with hydrogen and helium, the elements with the least mass, at the top of the table.
Figure 3.3 e Periodic Table of the Elements
(A) e elements are organized in the periodic table by atomic number. (B) Each element is found within its own box. e atomic number, atomic weight, name, and symbol are the most impo ant pieces of information.
In their most common form, many elements also contain the same number of neutrons as protons. e most common form of carbon, for example, has 6 neutrons as well as 6 protons, for a total of 12 subatomic pa icles in its nucleus. Since an atom’s atomic number is its number of protons, and its weight is roughly equal to the number of protons and neutrons, you can see in Figure 3.3B that atomic weight is approximately twice the value of the atomic number. e number of protons and electrons in an element are equal. e numbers of protons and neutrons may be equal for some elements but are not equal for all.
Isotopes
Each element has a unique number of protons, and typically the number of protons and neutrons are the same. However, some elements can have unequal numbers of protons and neutrons. ese variations are known as isotopes. An isotope is one of the various forms of an element, distinguished from one another by di erent numbers of neutrons. Examine the isotopes of hydrogen illustrated in Figure 3.4. All three of these atoms have one proton in their nucleus (illustrated in dark pink) and one electron in orbit (illustrated in light purple). Deuterium is the isotope of hydrogen that follows the standard format of an equal number of neutrons (illustrated in light yellow) to the protons and electrons. Protium is the most abundant isotope of hydrogen in nature; it has one proton and one electron but zero neutrons. Another isotope of hydrogen has two neutrons, but still one proton.
Figure 3.4Isotopes of Hydrogen
it donates one electron to another atom. After losing the electron, its outermost shell will have eight electrons, a much more stable con rmation. e loss will cause the positive charge of potassium’s protons to be more in uential than the negative charge of potassium’s electrons. In other words, the resulting potassium ion will be slightly positive. A potassium ion is written as , indicating that it has lost a single electron. A positively charged ion is known as a cation (Figure 3.5A).
Figure 3.5Potassium and Fluorine
(A) Potassium has 19 electrons, with just one lonely electron in its outermost shell. Potassium is likely to engage in reactions in which it donates that electron to another atom. In its new state, potassium has 18 electrons, a more stable 8 - electron outer shell, and a net positive charge. (B) Fluorine has nine electrons, with seven electrons in its outermost shell. Fluorine is likely to engage in reactions in which it accepts an electron from another atom to ll up its outermost shell. It now has a net negative charge because it has one more electron than it has protons.
Now consider uorine [F], a component of bones and teeth. Its atomic number is 9, and it has seven electrons in its valence shell. us, it is highly likely to bond with other atoms in such a way that uorine accepts one electron from another atom, lling its outermost shell to eight electrons, a more stable con rmation. When it does, its electrons will outnumber its protons by one, and it will have an overall negative charge. e ionized form of uorine is called uoride and is written as. A negatively charged ion is known as an anion (Figure 3.5B). Note that some ions get a nickname, while others do not. Cations, such as potassium, are simply referred to as potassium ions , but anions such as uorine and chlorine become uoride and chloride.
Student Study Tip
Remember that cations are positively charged because the word contains a “t,” just like a plus sign, for positive. “Anion” has 2 ns for negative.
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 as or. An anion that has accepted two electrons has a net charge of −2. e ionic form of selenium [Se], for example, is typically written as.
3.1c e Behavior of Electrons
In the human body, atoms do not hang out independently. Rather, they are constantly reacting with other atoms to form more complex substances. To fully understand anatomy and physiology you must grasp how atoms engage with each other. e key is in 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 ce ain regions of space called electron shells. An electron shell is a layer of electrons that encircle the nucleus at a distinct distance.
e atoms of the elements found in the human body have from one to ve electron shells, and all electron shells hold eight electrons except the rst shell, which can only hold two. is con guration of electron shells is the same for all atoms. e number of shells of the atoms of an element depends on the total number of electrons it has. Hydrogen and helium have just one and two electrons, respectively, and therefore these atoms have only one electron shell (Figure 3.6A). A second shell is necessa to hold the electrons in all elements larger than hydrogen and helium such as carbon (Figure 3.6B). Atoms with more than 10 electrons (such as sodium) require more than two shells (Figure 3.6C). Valence shell is the term for an atom’s outermost electron shell.
Figure 3.6Electron Shells Electrons orbit the atomic nucleus at distinct levels of energy called electron shells. (A) With one electron, hydrogen only half- lls its electron shell. Helium also has a single shell, but its two electrons completely ll it. (B) e electrons of carbon completely ll its rst electron shell, but only half- lls its second. (C) Sodium, with 11 electrons, lls both of its rst 2 electron shells and has just 1 electron in its valence shell.
e factor that most strongly governs the tendency of an atom to pa icipate in chemical reactions is the number of electrons in its valence shell. All atoms are most stable when their valence shell is happily full. 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- lling its valence shell. is 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.
In another 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, so that oxygen lls its valence shell and covalent bonds are formed, resulting in a molecule of water,.
e opposite charges of cations and anions attract the atoms into such close proximity that they form an ionic bond. An ionic bond is an ongoing, close association between ions of opposite charge. e most familiar example of ionic bonding is the table salt you sprinkle on your food. As shown in Figure 3.8, sodium commonly donates an electron to chlorine, becoming the cation. When chlorine accepts the electron, it becomes the chloride anion,. With their opposing charges, these two ions strongly attract each other forming NaCl, or salt.
Figure 3.8Ionic Bonding
(A) Sodium readily donates the solita electron in its valence shell to chlorine, which needs only one electron to have a full valence shell. (B) e 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) e attraction of many sodium and chloride ions results in the formation of large groupings called c stals.
3.2bCovalent 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 rst at one home and then at the other, the atoms do not lose or gain electrons
permanently. Instead, the electrons move back and fo h between the atoms. Because of the close sharing of pairs of electrons (one electron from each of two atoms), covalent bonds are stronger than ionic bonds.
Nonpolar and Polar Covalent Bonds
Figure 3.9 shows several common types of covalent bonds. Notice that the two covalently bonded atoms typically share one or two electron pairs. e impo ant concept to take from this is that in covalent bonds, electrons are shared to ll the valence shells of both atoms, ultimately stabilizing both of the atoms involved. In a single covalent bond, a single pair of electrons is shared between two atoms, while in a double covalent bond, two pairs of electrons (two electrons from one atom and two electrons from the other) are shared between two atoms. ere even are triple covalent bonds, where three atoms are shared.
Figure 3.9Covalent Bonding
(A) A single covalent bond joins two atoms of hydrogen. (B) A double covalent bond joins two atoms of oxygen. An atom of oxygen has six electrons in its valence shell; therefore, two more would make it stable. Two atoms of oxygen achieve stability by sharing two pairs of electrons in a double covalent bond. (C) Two double covalent bonds allow an atom of carbon, with four electrons in its valence shell, to achieve stability by sharing two pairs of electrons each with two atoms of oxygen.
You can see that the covalent bonds shown in Figure 3.9 are balanced. e sharing of the negative electrons is relatively equal, as they are anchored between the atoms by the positive pull of the protons on either side. Let’s contrast that with the water molecule illustrated in Figure 3.10. e water molecule has three pa s: one atom of oxygen, the nucleus of which contains eight protons, and two hydrogen atoms, each hydrogen nucleus containing only one proton. Because eve proton exe s an identical positive charge, a nucleus that contains eight protons exe s a charge eight times greater than a nucleus that contains one proton. is means that the negatively charged electrons present in the water molecule are more strongly attracted to the oxygen nucleus than to the hydrogen nuclei.
molecules simultaneously can be quite strong. Here a Green Basilisk or “Jesus Christ” lizard walks quickly across the su ace of water in a pond due to the lizard’s light weight and the cohesive prope y of the water.
e most common example of hydrogen bonding in the natural world occurs between molecules of water. It happens before your eyes whenever two raindrops merge into a larger bead, such as the beads of water seen in Figure 3.11B. Hydrogen bonds are individually weak and easily broken. We usually denote that weakness with a dotted line when we illustrate them, as opposed to the solid line used to illustrate the much stronger covalent bond. While they are weak individually, hydrogen bonds are strong collectively. Figure 3.11C is a photo of a small lizard that is capable of walking across the su ace of water, rather than falling in, held up by the cohesion among the water molecules.
Polar water molecules also strongly attract other polar molecules in the same way that they are attracted to each other (Figure 3.12A). Water molecules are also attracted to charged molecules and the charge on ions (Figure 3.12B). is explains why NaCl (table salt) dissolves so readily in water. Water readily dissolves all so s of polar or charged molecules, forming solutions. Water cannot, however, form hydrogen bonds with nonpolar molecules. In fact, water molecules repel molecules with nonpolar covalent bonds, like fats, lipids, and oils. You can demonstrate this with a simple kitchen experiment: pour a teaspoon of vegetable oil, a compound formed by nonpolar covalent bonds, into a glass of water. Instead of instantly dissolving in the water, the oil forms a distinct bead because the polar water molecules repel the nonpolar oil.
Figure 3.
(A) Water forms hydrogen bonds with other polar molecules in much the same way that it does with other water molecules. e pa ially positive charge of a hydrogen atom within water is attracted to the pa ially negative charge of (in this case) nitrogen within a polar molecule of ammonia. (B) e pa ially positive and negative ends of water are also attracted to anions and cations, respectively.
Hydrogen bonds are critical to the functioning of our molecules and cells. However, they are exquisitely vulnerable to environmental changes. Hydrogen bonds are likely to fall apa in the presence of acids or during changes in temperature.
3.3Chemical Reactions
One characteristic of a living organism is metabolism, which is the sum total of all of the chemical reactions that go on inside the body. e bonding processes you have learned thus far are key components in metabolism. Bonds are formed in order to build new molecules, and bonds are broken in the process of breaking down molecules. Anabolism is the process for building new molecules, and catabolism is the process for breaking molecules down.
3.3a e Role of Energy in Chemical Reactions
First, when thinking about atoms and molecules interacting, it is useful to de ne the di erent types of molecular energy. Energy is the capacity to do work, and the work of cells is to maintain order and structure as well as produce molecules such as hormones. e forms of energy relevant to anatomy and physiology are summarized in Table 3.1.
Table 3.1 Forms of Energy
Potential energy is stored energy that can be released. Kinetic energy is the energy of motion. Once potential energy is released, it becomes kinetic energy. Chemical energy is a form of potential energy in which energy is stored in bonds between atoms and molecules. When those bonds are formed, chemical energy is invested, and when they break, chemical energy is released.
Type of Energy Description Illustration Potential energy Stored energy ready to be released
Chemical reactions that release more energy than they absorb are characterized as exergonic reactions. e catabolism of the foods in your breakfast is an example. In contrast, chemical reactions that absorb more energy than they release are endergonic reactions (Figure 3.13). ese reactions require energy input, and the resulting molecule stores chemical energy. If energy is neither created nor destroyed, where does the energy needed for endergonic reactions come from? In many cases, it comes from exergonic reactions.
Figure 3.13Endergonic and Exergonic Reactions
Endergonic reactions require energy; the energy is stored within bonds as potential energy. Exergonic reactions release energy through the breaking of bonds. e most common example of these reactions in A&P is the formation and breakdown of the energy storage molecule, ATP.
3.3bCharacteristics of Chemical Reactions
All chemical reactions begin with a reactant, the general term for the one or more substances that enter into the reaction. Sodium and chloride ions, for example, are the reactants in the production of table salt. e one or more substances produced by a chemical reaction are called the product. In the gure in the “Anatomy of Chemical Reactions” feature, A and B represent reactants and AB represents the product.
Student Study Tip
As shown here, practice how di erent types of chemical reactions run using variables, drawings, words, and so on to make sure you understand the mechanisms before adding molecules!
Just as you can express mathematical calculations in equations such as , you can use chemical equations to show how reactants become products. As in math, chemical equations proceed from left to right, but instead of an equals sign, they employ an arrow or arrows indicating the direction in which the chemical reaction proceeds. For example, the chemical reaction in which one atom of nitrogen and three atoms of hydrogen produce ammonia would be written as. Correspondingly, the breakdown of ammonia into its components would be written as.
Notice that, in the rst example, a nitrogen [N] atom and three hydrogen [H] atoms bond to form a compound. is reaction forms bonds, so we can describe it as anabolic and it requires energy, which is then stored within the compound’s bonds, so we can describe it as endergonic. Such reactions are referred to as synthesis reactions. A synthesis reaction is a chemical reaction that results in the synthesis (joining) of components that were formerly separate (pa A in the “Anatomy of Chemical Reactions” feature). Again, nitrogen and hydrogen are reactants in a synthesis reaction that yields ammonia as the product. e general equation for a synthesis reaction is.
In the second example (pa B of the “Anatomy of Chemical Reactions” feature), ammonia is catabolized into its smaller components, and the potential energy that had been stored in its bonds is released. Such reactions are referred to as decomposition reactions. A decomposition reaction is a chemical reaction that breaks down a molecule into its constituent pa s. e general equation for a decomposition reaction is:.
e third, and the most complex reaction in pa C of the “Anatomy of Chemical Reactions” feature is an exchange reaction, a chemical reaction in which both synthesis and decomposition occur, chemical bonds are both formed and broken, and chemical energy is absorbed, stored, and released. e exchange reaction shown can be summarized as:. Notice that AB and CD had to be broken down in decomposition reactions to produce these products, whereas A and C and B and D had to bond in synthesis reactions.
Anatomy of… Chemical Reactions
Chemical reactions in the body can be broken into three types: synthesis, decomposition, and exchange reactions.
In theo , any chemical reaction can proceed in either direction under the right conditions. Reactants may synthesize into a product that is later decomposed. is reversibility of a chemical reaction is indicated with a double arrow:. Still, in the human body, many chemical reactions do proceed in a predictable direction, either one way or the other. You can think of this more predictable path as the path of least resistance because, typically, the alternate direction requires more energy. In the wate environment of the human body, we tend to see primarily ce ain types of reactions that occur in water (dehydration synthesis and hydrolysis reactions).
3.3cFactors In uencing the Rate of Chemical Reactions
If you pour vinegar into baking soda, the reaction is instantaneous; the concoction will bubble and zz. But many chemical reactions take time. A variety of factors in uence the rate of chemical reactions. is section, however, will consider only the most impo ant factors in the reactions that take place within human bodies and cells. e most impo ant thing to remember about the likelihood of a synthesis reaction to occur is that the reactants need to be close together in order to interact. at may seem obvious, but remember that these reactants are in nitely small
Student Study Tip
Activation energy cu es act like a ball being rolled up a hill: once over the hump, the ball will roll down without e o . A smaller hump leads to quicker/easier reactions/rolling; sho ening the hump height explains enzyme function.
e most impo ant catalysts in the human body are enzymes. An enzyme is a protein that makes chemical reactions faster and more likely to occur by lowering activation energy. For example, if the activation energy in a synthesis reaction is due to two molecules being close enough together for long enough to form bonds, then an enzyme lowers the activation energy by holding them close together as illustrated in Figure 3.14B. Enzymes are critical to the body’s healthy functioning. ey assist, for example, with the breakdown of food and its conversion to energy and the building of many cellular molecules. In fact, most of the chemical reactions in the body are facilitated by enzymes.
3.4Inorganic Compounds Essential to Human Functioning
e next two sections of the chapter cover the compounds impo ant for the body’s structure and function. In general, these compounds are either inorganic or organic. Table 3.2 lists some of the common molecules and compounds referenced throughout this book.
Table 3.2 Commonly Used Chemical Formulas in A&P
Full Name Chemical Formula
e following section examines the three groups of inorganic compounds essential to life: water, salts, and acids and bases. Organic compounds are covered later in the chapter.
3.4aWater
Scientists estimate that between 50 and 70 percent of an adult’s body is composed of water (this number varies pretty widely and is in uenced by age, body composition, and the levels of hormones, pa icularly estrogen). is water is contained both within the cells and between the cells that make up tissues and organs. Its several roles make water indispensable to human functioning.
Water as a Lubricant and Cushion
Water is a major component of many of the body’s lubricating uids. Just as oil lubricates the hinge on a door, water in our joints lubricates the tissues of those joints (explored in Chapter 10 ), and water inside of cells is critical for their function (explored in Chapter 4 ). Wate uids help keep food owing through the digestive tract and ensure that the movement of adjacent abdominal organs is friction free (Chapter 24 ).
Water also protects cells and organs from physical trauma, cushioning the brain within the skull, for example, and protecting the delicate ne e tissue of the eyes. Water cushions a developing fetus within the uterus as well. Water is fundamental to human life and plays many di erent roles in the human body:
Water as a Heat Sink
A heat sink is a substance or object that absorbs and dissipates heat but does not experience a corresponding increase in temperature. In the body, water absorbs the heat generated by chemical reactions without greatly increasing in temperature. Moreover, when the environmental temperature soars, the water stored in the body helps keep the body cool. Water brings heat with it as it changes from a liquid to a gas phase. is mechanism, known as evaporative cooling , explains how sweat helps the body to cool down once it has warmed past homeostatic range.
Oxygen Carbon dioxide Potassium ions Chloride ions Calcium ions (also expressed ) Glucose Water Hydrochloric acid HCl Sodium chloride NaCl
An inorganic compound is a substance that does not contain both carbon and hydrogen. A great many inorganic compounds do contain hydrogen atoms, such as water and the hydrochloric acid (HCl) produced by your stomach. In contrast, only a handful of inorganic compounds contain carbon atoms. Carbon dioxide is one of the few examples. An organic compound, then, is a substance that contains both carbon and hydrogen. Organic compounds are synthesized via covalent bonds within living organisms, including the human body.
Water is the main component of blood and lymph, and se es a transpo ation role, moving cells, nutrients, hormones, and wastes around the body. rough sweating, water helps to bring elevated body temperatures down. A wate uid acts as a cushion around the brain, protecting it from harm during sudden movements. Water is the solvent of the body for hydrophilic molecules. e behavior of hydrophobic molecules is de ned by their interactions with the water. In our joints and around our organs, water-based uids lubricate and reduce friction.
