Eia ka wai la, he wai e ola. E ola nō e!

Here is the water, the water of life. Life, indeed!

He Pule Kuahu (no Kini Akua)

Introduction

Figure 2.1. Some Hawaiian Dishes. From top left, clockwise: tripe stew, rice, opihi poke, laulau, squid luau, pipikaula short ribs, kalua pig, and poi in the center. (Credit: By takaokun, CC BY 2.0, https://commons.wikimedia.org/w/index.php?curid=44626828)

Imagine you are at a potluck with your family and friends. You have all your favorite foods such as laulau, short ribs, poi, and rice. All of these foods are types of matter. They contain flavor molecules, ions, salt, water, and macromolecules such as proteins and carbohydrates. Have you ever wondered why we need these substances, and what they do in our bodies?

In this chapter, we will learn about the chemical structures of these substances and how they contribute to the overall function of the body.

2.1 Matters, Elements, and Atoms

Matters

Your favorite laulau, short ribs, poi, and rice all consist of matter. Matter is any substance that has mass and occupies space. An object’s mass is the amount of matter it contains. 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. For example, short ribs that weigh a pound on Earth weigh only a few ounces on the moon, even though the mass of the ribs remains the same.

Chemical Elements

All matter in the natural world is composed of one or more of the 118 fundamental substances called elements. Elements such as pure gold are pure substances that cannot be broken down into simpler substances by chemical changes. Each element is designated by its chemical symbol, which is typically a single capital letter or a combination of two letters.

A familiar example of an element in the human body is calcium and its chemical symbol is Ca. Calcium is absorbed and used for many 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. The elemental calcium in cheese and milk, therefore, is the same as the calcium that forms your bones.

Common chemical elements in the human body:

Figure 2.2 Elements of the Human Body The main elements that compose the body are shown from most abundant to least abundant.

In the human body, 26 elements are derived from the foods you eat and the air you breathe [Figure 2.2]. The four most abundant elements in the body are oxygen (O), carbon (C), hydrogen (H), and nitrogen (N), and together they make up 96% of the body’s mass.

In nature, elements rarely occur alone. Instead, two or more elements are joined by chemical bonds to form compounds such as glucose, which is an important body fuel. Glucose is always composed of the same three elements: carbon, hydrogen, and oxygen, and this always occurs in the same relative amounts (e.g. six carbon and six oxygen units for every twelve hydrogen units as in C₆H₁₂O₆. But what, exactly, are these “units” of elements?).

Structure of an Atom

An atom is the smallest quantity of an element that still has 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.

Proton, Neutron, Electron

Atoms are made up of even smaller subatomic particles, three types of which are important: proton, neutron, and electron. Protons and neutrons are larger subatomic particles found in the nucleus, which is in the center of the atom. Compared to the other two subatomic particles, electrons are much smaller — about 1/2000th the mass of a proton or neutron. Electrons orbit (“spin”) around the nucleus, similar to planets circling the sun.

(c) An animated atomic model.

Figure 2.3 Models of Atomic Structure (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. Note that the electrons in reality are much smaller than the protons and neutrons depicted here. (b) In the electron cloud model, the electrons of helium are shown in the variety of locations they would have at different distances from the nucleus over time. The electron cloud model more accurately shows the relative structure of an atom. (c) An animated representation of the atomic model with electrons in yellow orbiting the nucleus. (By LG UltraLink – Own work, CC BY-SA 4.0).

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 and are electrically neutral. This means that even though protons and neutrons exist in the nucleus of an atom, the nucleus has a positive charge as seen in Figure 2.3b. 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. For example, a helium atom in Figure 2.3c has two protons and two electrons, making this a neutral atom.

Electron Shells and Valence Shell

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 encircles the nucleus at a distinct energy level. For example, in Figure 2.3c, two electrons in the helium atom are shown circling the nucleus in a fixed orbit depicted as a ring. Although this model helps visualize the atomic structure, in reality, electrons do not travel in fixed orbits, but whiz around the nucleus erratically in a so-called electron cloud (Figure 2.3b).

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 (Figure 2.4). 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, and so only have one electron shell as seen in Figure 2.4a. Additional shells are necessary to hold the electrons in all elements larger than hydrogen and helium.

Figure 2.4 Electron Shells 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 fill it. (b) The electrons of carbon fill its first electron shell, but only half-fill its second. (c) Neon, an element that does not occur in the body, has 10 electrons, filling both of its electron shells.

Carbon (C), with its six electrons, fills its first shell and half-fills its second (Figure 2.4b). With ten electrons, neon (Ne) fills its two electron shells (Figure 2.4c). Atoms with more than ten electrons require more than two shells.

A valence shell is an atom’s outermost electron shell. 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/s so that it ends up with eight electrons in its own valence shell. In other words, the atom isn’t likely to engage in chemistry until its orbitals are filled. For example, carbon, with four electrons in its valence shell, is likely to react with other atoms in a way that results in the addition of four electrons to carbon’s valence shell, bringing the number to eight.

In nature, atoms of one element tend to join with atoms of other elements in characteristic ways. One example is water; oxygen needs two electrons to fill its valence shell. It commonly interacts with two atoms of hydrogen, forming H2O. Incidentally, the name “hydrogen” reflects its contribution to water (“hydro-” = “water”; “-gen” = “maker”). Thus, hydrogen is the “water maker.”

Atomic Number and Mass Number

If all atoms are made of the same three subatomic particles, what makes a carbon atom so different from a sodium or iron atom? The answer is the unique quantity of protons each contains. Carbon by definition is the only element whose atoms contain six protons (see Figure 2.5). No other element has exactly six protons in its atoms. Moreover, all atoms of carbon, whether found in your genes or proteins, contain six protons. Thus, the atomic number, which is the number of protons in the nucleus of the atom, identifies the element. Each atom has a specific number of protons. Because an atom usually has the same number of electrons as protons, the atomic number identifies the usual number of electrons as well. To reiterate, the atomic number (the number at the top of the element in the periodic table) represents the number of protons (Figure 2.6). In addition, it is also representative of BOTH the number of protons and electrons. If the atom does not contain an overall charge it can be assumed that the number of protons and electrons are equal for a neutral atom.

Figure 2.5 The atomic number and mass number of carbon twelve (C-12). (https://chem.libretexts.org/Courses/Howard_University/General_Chemistry%3A_An_Atoms_First_Approach/Unit_1%3A__Atomic_Structure/Chapter_1%3A_Introduction/Chapter_1.6%3A_Isotopes_and_Atomic_Masses).

In their most common form, many elements also contain the same number of neutrons as protons. An element’s mass number is the number of protons and neutrons in its nucleus. For carbon, the most common form has six neutrons and six protons which makes its mass number equal to 12. The mass number for carbon is 12.01 based on the periodic table because it is an average mass number of all existing types of carbons (Figure 2.6). Electrons have so little mass that they do not appreciably contribute to the mass of an atom. The number of protons and neutrons may be equal for some elements but are not equal for all. For example, iodine (I) has a mass number of 127. Its atomic number is 53 (it has 53 protons) but it contains 74 neutrons.

Periodic Table

The periodic table of the elements is a chart identifying the 118 elements discovered or synthesized so far (Figure 2.6). 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.

Figure 2.6 The Periodic Table of the Elements (credit: R.A. Dragoset, A. Musgrove, C.W. Clark, W.C. Martin)

Isotopes

Isotopes are different forms of an element that have the same number of protons but a different number of neutrons. Some elements — such as carbon, potassium, and phosphate, have naturally occurring isotopes. The standard isotope of carbon is carbon-12 (12C), commonly called carbon twelve which has six protons and six neutrons. All of the isotopes of carbon have the same number of protons; therefore, 13C has seven neutrons, and 14C has eight neutrons. The different isotopes of an element can also be indicated with the mass number hyphenated (for example, C-12 instead of 12C).

Figure 2.7 Isotopes of Carbon. The isotopes of carbon have the same number of protons but different numbers of neutrons.

Radioisotope

Certain isotopes are radioisotopes such as 14C. It contains more than the usual number of neutrons and tends to be unstable. When 14C breaks down and becomes more stable, it emits radiation. Excessive exposure to radioisotopes can damage human cells and even cause cancer and birth defects, but when exposure is controlled, some radioactive isotopes can be useful in medicine. The controlled use of radioisotopes has advanced medical diagnosis and treatment of disease. One of the most advanced uses of radioisotopes in medicine is the positron emission tomography (PET) scanner. The PET scanner detects the activity of a very small amount of radioactive glucose, the simple sugar that cells use for energy, injected into the body (Figure 2.8). The PET camera captures which of the patient’s tissues are taking up the most glucose. Thus, the most metabolically active tissues (those that are using the most glucose) show up as bright “hot spots” on the images. PET can reveal some cancerous masses because cancer cells consume glucose at a higher rate than healthy cells to fuel their rapid reproduction.

Figure 2.8 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.

Molecules and Compounds

Atoms interact with each other by forming a chemical bond. A bond is an electrical attraction that can have varying degrees of strength holding 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 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 H2, which is called molecular hydrogen or hydrogen gas. A molecule can also be formed from different elements, such as H2O. 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 H2O, is also a compound, as is a single molecule of the gas methane, or CH4.

Free Radicals and Antioxidants

Free radicals are molecules and ions with an odd number of electrons in their valence shells. These unpaired electrons make them highly chemically reactive. In the human body, free radicals can form from natural metabolism of oxygen when it accepts the unpaired electrons. For example, when oxygen gas is bonded to one unpaired electron, it forms superoxide (•O−2).  Although free radicals play a role in several physiological functions in the human body, free radicals can also react with biological molecules such as DNA, RNA, proteins, and fatty acids and cause damage to cells. For example, a superoxide molecule may react with a fatty acid and disrupt the cell membrane or proteins in cell membranes. Substances in the food you eat, such as fried foods, or drinks, such as alcohol, can generate free radicals.

The human body can also be exposed to free radicals from a variety of environmental sources, such as cigarette smoke, air pollution, fried food, and sunlight. Highly reactive free radicals are generated when using an electronic cigarette (used when vaping). Free radicals can cause “oxidative stress,” a process that can trigger cell damage. Oxidative stress is thought to play a role in a variety of diseases including cancer, cardiovascular diseases, diabetes, Alzheimer’s disease, Parkinson’s disease, and eye diseases such as cataracts and age-related macular degeneration. Antioxidants are chemicals largely responsible for protecting cell membranes from free radical damage. Some antioxidants are synthesized by the human body and others are found in the food we eat, including ‘olena (turmeric), avocados, and sweet potatoes.

2.2 Chemical Bonds

When you are eating your favorite salad, have you ever wondered about what is present in your salad dressing? In a salad dressing, there is usually the presence of salt, water, and other molecules to enhance the flavor. In this section, we will learn about ions, water, and other molecules as well as the chemical bonds that hold them together. Remember that all matter is made from atoms, so these atoms must come close enough for the electrons in their valence shells to interact. Interaction between electrons in an atom’s valence shells with those of another atom can result in a chemical bond.

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.

Ionic Bonds (Cations, Anions, and Electrolytes)

An atom that has an electrical charge — whether positive or negative — is an ion. The common salt in a salad dressing is composed of two ions, Na+ and Cl- (see Figure 2.9c). These two ions have opposite charges and are held together by ionic bond forces. Let’s take a look at how these ions are formed. When an atom has the same number of positively charged protons and negatively charged electrons, the atom is electrically neutral. For example, a neutral sodium atom (Na) contains 11 protons and 11 electrons, whereas a neutral chlorine atom (Cl) contains 17 protons and 17 electrons.

However, just because an atom is neutral, it doesn’t mean that the neutral atom is stable. To become stable, an atom will try to fill its valence shell. This can happen either by gaining electrons to fill an almost full shell or by giving away electrons from a nearly empty shell. For example, a sodium atom (Na) has only one electron in its valence shell so it loses the electron from its valence shell to form a more stable sodium ion (Figure 2.9a). A sodium ion is written as Na+, indicating that it has lost a single negatively charged electron, which results in a net positive charge as there is now one more proton than there are electrons. A positively charged ion is known as a cation.

On the other hand, a chlorine atom (Cl) has seven electrons in its valence shell and by accepting one electron, it fills its valence shell to form a stable chloride ion (Figure 2.9a). When it does this, its electrons will outnumber its protons by one, and it will have an overall negative charge. The ionized form of chlorine is called chloride, and is written as Cl–. A negatively charged ion is known as an anion. This happens frequently for most atoms to have a full valence shell, as described previously.

Figure 2.9 Ionic Bonding (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.

Some 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 calcium (Ca) as an example, this can be written Ca++ or Ca2+. An anion that has accepted two electrons has a net charge of –2. The ionic form of selenium (Se), for example, is typically written Se2–.

The opposite charges of cations and anions exert a mutual attraction that keeps the atoms in close proximity forming an ionic bond. The crystals of table salt that you sprinkle on your food owe their existence to ionic bonds that position the sodium and chloride ions into an alternating pattern in a matrix (Figure 2.9b).

Role of Electrolytes in the Human Body

It takes a tremendous amount of energy to break the ionic bond within salt. For example, if you drop a large block of salt, you tend to just get smaller blocks of salt. However, a typical salt, such as NaCl, dissociates (separates) completely in water without the changing of temperature (Figure 2.10). Have you wondered why water can do this so easily? Take a look at Figure 2.9 and you will see that water can break the ionic bonds in salts by surrounding each individual ion. For example, the positive and negative regions on the water molecule (the hydrogen and oxygen ends, respectively) attract the negative chloride and positive sodium ions, pulling them away from each other. This explains why “table salt” (NaCl) which consists of equal numbers of sodium ions (Na+) and chloride ions (Cl–), dissolves so readily in water, in this case forming dipole-ion bonds between the water (i.e. dipole) and the electrically-charged ions (electrolytes).

Figure 2.10 Dissociation of Sodium Chloride in Water Notice that the crystals of sodium chloride dissociate into Na+ cations and Cl– anions, each surrounded by water molecules.

In biological fluids, many ionic compounds exist as ions as they are dissolved by water in the body. In the human body, Na+, potassium (K+), Cl-, Ca2+ are examples of ions that play an important role in various physiological processes. For example, Na+ and K+ play a role in membrane potential that is important for carrying nerve impulses. Besides forming the solids of calcium phosphate in strengthening the bones and teeth, Ca2+ also has an important role in muscle contraction. These charged ions are called electrolytes because they produce electrical charges within the body. The function of the heart and brain act to influence the overall behavior of these ions and the resulting electrical activity can be detected and observed as waves on an electrocardiogram (ECG or EKG) or an electroencephalogram (EEG).

Homeostatic imbalances and clinical connections related to ions in the body

Although ions are essential for the human body, there are times when having too many or too few ions can lead to potential diseases in the long term. For example, having an imbalance of Na+ and K+ in the body can affect the conduction of nerve impulses and various physiological processes. Also, there is a strong correlation between high levels of sodium ions and hypertension (e.g. high blood pressure). Be mindful that hypertension can be caused by different factors.

Covalent Bonds

Unlike ionic bonds, molecules formed by a covalent bond share electrons in a mutually stabilizing relationship. Similar to two children who share their toys, these atoms share their electrons. The atoms do not lose or gain electrons permanently. Because of the close sharing of pairs of electrons (one electron from each of two atoms), covalent bonds are stronger than ionic bonds in biological systems.

Nonpolar covalent bond

When you imagine two children who share their toys well, the toys are mostly shared equally. When two atoms share their electrons equally, they form a nonpolar covalent bond which means that the sharing of electrons between both atoms is fairly well balanced. Figure 2.11 shows several common types of nonpolar covalent bonds. The important concept to take from this is that with covalent bonds, the two atoms pull the electrons equally, thus the naming of “nonpolar” where no region of the molecule is either more positive or more negative than any other. The electrons are shared to fill the valence shells of both atoms, ultimately stabilizing both of the atoms involved. In a single covalent bond, a single electron pair (i.e. 2 electrons) is shared between two atoms (Figure 2.11a), while in a double covalent bond, two pairs of electrons (i.e. 4 electrons) are shared between two atoms such as between two oxygen atoms forming the oxygen gas that we breathe in (Figure 2.11b and c). There even are triple covalent bonds, where three electron pairs (i.e. 6 electrons) are shared between two atoms.

Figure 2.11 Covalent Bonding Different types of nonpolar covalent bonds are shown here where the atoms share electrons equally in a single or double covalent bond.

Polar Covalent Bond

Let’s go back to the toy analogy. Imagine two children are trying to share one toy, but the toy is now being pulled towards one child more than the other. In this situation, the toy is being shared, but not equally between the two children. When two atoms share electrons unequally, the atom that has its electron pulled away from it ends up with a partial positive charge and the other atom that more forcefully pulls the electron toward it will assume a partial negative charge. This type of electron sharing is known as a polar covalent bond (see Figure 2.12).

In chemistry, a polar molecule is a molecule that contains regions that have opposite electrical charges due to polar covalent bonds. The most familiar example of a polar molecule is water (Figure 2.12). The water molecule has three parts: one oxygen atom and two hydrogen atoms. The nucleus of an oxygen atom contains eight protons, while nuclei of hydrogen atoms each contain only one proton. Because every proton exerts an identical positive charge, an oxygen nucleus that contains eight protons exerts a charge eight times greater than a hydrogen nucleus that contains one proton. This means that the negatively charged electrons present in the water molecule are more strongly attracted to the oxygen nucleus than to the hydrogen nuclei. Therefore, a negatively charged electron from each hydrogen atom moves towards the oxygen atom, making the oxygen end of their bond slightly more negative than the hydrogen end of their bond. These charges are often referred to as “partial charges” because the strength of the charge is less than one full electron, unlike in an ionic bond. Thus, we say that the water molecule is “polar”. It has a partial negative charge in the oxygen region and a partial positive charge in the regions of the hydrogen atoms.

As shown in Figure 2.12, regions of partial positive or negative charges in a polar molecule are indicated with the Greek letter delta (δ) and a plus (+) or minus (–) sign.

Figure 2.12 Polar Covalent Bonds in a Water Molecule The oxygen atom is partially negatively charged while the hydrogen atoms are partially positively charged.

Electronegativity

So how do you predict if two atoms will share electrons equally (producing a nonpolar covalent bond) or share them unequally (producing a polar covalent bond)?

This is determined by a property of the bonding atoms called electronegativity. Electronegativity is the tendency of an atom to attract electrons (or electron density) towards itself. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. In a polar covalent bond, electrons are shifted toward the more electronegative atom of the two bonding atoms; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity between the two bonding atoms, the more polarized the electron distribution (i.e. electrons distribute closer to the more electronegative atom) and results in the larger the partial charges of the atoms. In the biological systems, we will see cases where oxygen and nitrogen atoms being more electronegative than hydrogen and carbon atoms. For instance, water is an example where the electronegative oxygen pulls electrons from the two hydrogen atoms (see Figure 2.12).

Hydrogen Bonds

The third and last type of chemical bond that we will discuss is a hydrogen bond. In this section, you will find that hydrogen bonds are a force that holds water molecules, DNA, and protein together.

H-bonds Between Water Molecules

The 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. Hydrogen bonding occurs because the partially negative oxygen atom in one water molecule is attracted to the partially positive hydrogen atoms of two other water molecules (Figure 2.13).

Figure 2.13 Hydrogen Bonds between Water Molecules Notice that the bonds occur between the weakly positive charge on the hydrogen atoms and the weakly negative charge on the oxygen atoms. Hydrogen bonds are relatively weak and therefore are indicated with a dotted (rather than a solid) line.

Hydrogen bonds between liquid water molecules are the forces that we try to break with heat when we boil a pot of water to make coffee. This gives rise to a special property of water which will be discussed in the next section. Collectively, the hydrogen bonds between water molecules are strong enough for an insect known as a water strider to walk on water without sinking (see Figure 2.14).

Figure 2.14 Image of a water strider walking on water (By © Nevit Dilmen, CC BY-SA 3.0, https://commons.wikimedia.org/w/index.php?curid=27930507)

A unique characteristic of polar water is that it repels molecules with nonpolar covalent bonds, such as 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.

H-bonds Between Nucleotides and Proteins

Hydrogen bonds also play an essential role in holding the DNA molecules together and creating the three-dimensional shape of proteins such as antibodies (Figure 2.15). In these cases, hydrogen bonds can form between a hydrogen atom of a polar molecule and a slightly negative atom (oxygen or nitrogen) of another polar molecule.  Without hydrogen bonds, DNA and antibodies would not be able to form their unique shapes. In human anatomy and physiology, the shape is intricately linked to function. Thus, for antibodies, a change in shape will affect their ability to bind to pathogens and trigger immune responses.

a

b

Fig 2.15 Hydrogen bonds between bases of DNA and proteins. (By CNX OpenStax — http://cnx.org/contents/GFy_h8cu@10.53:rZudN6XP@2/Introduction, CC BY 4.0, https://commons.wikimedia.org/w/index.php?curid=49922725) (By CNX OpenStax — http://cnx.org/contents/GFy_h8cu@10.53:rZudN6XP@2/Introduction, CC BY 4.0, https://commons.wikimedia.org/w/index.php?curid=49923721)

2.3 Chemical Reactions

When you eat rice, the chemical bonds that hold large molecules such as starch are broken down by the components of the digestive system, releasing smaller molecules. The chemical reactions that break the bonds between components of larger molecules are called catabolic reactions. On the other hand, chemical reactions that form bonds between small molecules to create larger ones are called anabolic reactions. Metabolism is the sum of all chemical reactions, both catabolic and anabolic, that maintain an organism’s health and life. Both types of chemical reactions involve exchanges, not only of matter but of energy, to break existing chemical bonds and form new ones.

Forms of Energy and Law of Conservation of Energy

Kinetic, Potential, and Chemical Energy

Kinetic energy is the energy of motion, meaning that a form of energy powering any type of matter in motion, such as swimming or running. Potential energy is the energy of position, or the energy matter possesses because of the positioning or structure of its components. In the body, potential energy is stored in the chemical bonds between atoms and molecules. This type of potential energy is called chemical energy. When those bonds are formed, chemical energy is invested, and when the bonds break, chemical energy is released. Notice that chemical energy, as with all energy, is neither created nor destroyed; rather, it is converted from one form to another. This describes the Law of Conservation of Energy. When you eat an energy bar before heading out the door for a swim or a run, the honey, nuts, and other ingredients in the bar are broken down and rearranged by your body into smaller molecules that your muscle cells can convert to kinetic energy.

How does this apply to body function?

Chemical reactions that release more energy than they absorb are referred to as exergonic reactions (energy EXITs). We can apply this in the energy bar example. Some of the chemical energy stored in the bar is absorbed into molecules your body uses for fuel, but some of it is released — for example, as heat. In contrast, chemical reactions that absorb more energy than they release are endergonic reactions. These reactions require energy input, and the resulting molecule stores not only the chemical energy in the original components but also the energy that fueled the reaction. Because energy is neither created nor destroyed, where does the energy needed for endergonic reactions come from? In many cases, it comes from exergonic reactions.

You have already learned that chemical energy is absorbed, stored, and released by chemical bonds. In addition to chemical energy, mechanical, radiant, and electrical energy are important in human functioning.

Mechanical energy is the sum of kinetic energy and potential energy. Mechanical energy, which is stored in physical systems such as machines, engines, or the human body, directly powers the movement of matter. When you lift a brick into place on a wall, your muscles provide the mechanical energy that moves the brick.

Radiant energy is energy emitted and transmitted as waves rather than matter. These waves vary in length from long radio waves and microwaves to short gamma waves emitted from radioisotopes. The full spectrum of radiant energy is referred to as the electromagnetic spectrum. The body uses the ultraviolet energy of sunlight to convert a compound in skin cells to vitamin D, which is essential to human functioning. The human eye evolved to see the wavelengths that comprise the colors of the rainbow, from red to violet, so that range in the spectrum is called “visible light.”

Electrical energy, supplied by electrolytes in cells and body fluids, contributes to the voltage changes that help transmit impulses in nerve and muscle cells.

Characteristics of Chemical Reactions

Chemical Reactions Include Reactants and Products

All chemical reactions begin with a reactant, the general term for one or more substances that enter into the reaction. One or more substances produced by a chemical reaction are called the product. Sodium and chloride ions, for example, are the reactants in the production of table salt.

In chemical reactions, the components of the reactants — the elements involved and the number of atoms of each — are all present in the product(s) but in different arrangements. There should not be any new component(s) or loss of component(s) in the product(s). This is because chemical reactions are governed by the Law of Conservation of Mass, which states that matter cannot be created or destroyed in a chemical reaction.

Just as you can express mathematical calculations in equations such as 4 + 2 = 6, 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 equal sign, an arrow or arrows are used to indicate the direction in which the chemical reaction proceeds. For example, the chemical equation 2 H2 + O2 → 2 H2O, shows that two hydrogen molecules (H2) and one oxygen molecule (O2) are the reactants and two water molecules (H2O) are the products of this chemical reaction. When you count the total number of atoms in the reactants, there are four hydrogen and two oxygen atoms. Thus, in the products (two H2O molecules), the same number of hydrogen and oxygen atoms are present.

Synthesis, Decomposition, Exchange, and Reversible Reactions

Notice that, in the example above, hydrogen (H) atoms and oxygen (O) atoms bond to form a compound. This anabolic reaction requires energy, which is then stored within the compound’s bonds. Such reactions are referred to as synthesis reactions. A synthesis reaction is a chemical reaction that results in the synthesis (joining) of formerly separate components (Figure 2.16a). The general equation for a synthesis reaction is A + B→AB.

Figure 2.16 The Three Fundamental Chemical Reactions The atoms and molecules involved in the three fundamental chemical reactions can be imagined as words.

When a bond within a compound is broken and results in its smaller components, 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 or “decomposes” something larger into its constituent parts (see Figure 2.16b). The general equation for a decomposition reaction is: AB→A+B.

An exchange reaction is 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 (see Figure 2.16c). The simplest form of an exchange reaction might be: A+BC→AB+C. Notice that, to produce these products, B and C had to break apart in a decomposition reaction, whereas A and B had to bond in a synthesis reaction. A more complex exchange reaction might be: AB+CD→AC+BD. Another example might be: AB+CD→AD+BC.

In theory, any chemical reaction can proceed in either direction under the right conditions. Reactants may synthesize into a product that is later decomposed. Reversibility is also a quality of exchange reactions. For instance, A+BC→AB+C could then reverse to AB+C→A+BC. This reversibility of a chemical reaction is indicated with a double arrow: A+BC⇄AB+C. Still, in the human body, many chemical reactions do proceed in a predictable direction, either one way or the other — not both. You can think of this more predictable path as the path of least resistance because, typically, the alternate direction requires more energy.

Oxidation-reduction reactions

When one reactant loses electrons (“oxidized”) while another gains electrons (“reduced”) in a chemical reaction, this chemical reaction is considered an oxidation-reduction reaction, also called a redox reaction. For example, glucose is fuel for our cells and it undergoes a series of redox reactions through cellular respiration. The main product of these reactions is chemical energy for our cells to work. Details on redox reactions will be discussed in Chapter 24 which cover the metabolism.

Figure 2.17 Example of Redox Reaction Sodium loses its outer electron to give it a stable electron configuration, and this electron enters the fluorine atom. Sodium and fluorine bonding ionically to form sodium fluoride. (By Wdcf – Own work, CC BY-SA 3.0, https://commons.wikimedia.org/w/index.php?curid=15539913)

Factors Influencing the Rate of Chemical Reactions

If you pour vinegar into baking soda, the reaction is instantaneous; the concoction will bubble and fizz. Chemically, acetic acid and sodium bicarbonate react to form carbon dioxide (i.e. bubble), water, and sodium acetate. But many chemical reactions take time to be fully completed. Chemical reactions in the body can also be influenced by outside factors. A variety of factors influence the rate of chemical reactions. This section, however, will consider only the most important in human functioning.

Temperature, Concentration, and Pressure

Nearly all chemical reactions occur at a faster rate at higher temperatures. Recall that kinetic energy is the energy of matter in motion. The kinetic energy of subatomic particles increases in response to increases in thermal energy. The higher the temperature, the faster the particles move, and the more likely they are to come in contact and react to form or break a bond.

If just a few people are dancing at a club, they are unlikely to step on each other’s toes. But as more and more people get up to dance — especially if the music is fast — collisions are likely to occur. It is the same with chemical reactions: the more particles present within a given space, the more likely those particles are to bump into one another. This means that chemists can speed up chemical reactions not only by increasing the concentration of particles — the number of particles in the space — but also by decreasing the volume of the space, which would correspondingly increase the pressure. If there were 100 dancers in that club, and the manager abruptly moved the party to a room half the size, the concentration of the dancers would double in the new space, and the likelihood of collisions would increase accordingly.

Catalysts, Enzymes, and Activation Energy

For two reactants in nature to react with each other, they first have to come into contact, and this occurs through random collisions. Because heat helps increase the kinetic energy of atoms, ions, and molecules, it promotes their collision. But in the body, extremely high heat — such as a very high fever — can damage body cells and be life-threatening. On the other hand, normal body temperature is not high enough to promote the chemical reactions that sustain life. That is where catalysts come in.

In chemistry, a catalyst is a substance that increases the rate of a chemical reaction without itself undergoing any change. You can think of a catalyst as a chemical change agent which provides the environment to increase the rate and force at which atoms, ions, and molecules collide, thereby increasing the probability that the valence shell electrons will interact.

The most important catalysts in the human body are enzymes. An enzyme is a catalyst composed of protein or ribonucleic acid (RNA), both of which will be discussed later in this chapter. As with all catalysts, enzymes provide the environment for the reactants and thus it lowers the level of activation energy that needs to be invested in a chemical reaction (see Figure 2.18). A chemical reaction’s activation energy is the “threshold” level of energy needed to break the bonds in the reactants. Once those bonds are broken, new arrangements can form. Without an enzyme to act as a catalyst, a much larger investment of energy is needed to ignite a chemical reaction.

Enzymes are critical to the body’s healthy functioning. They assist, for example, with the breakdown of food and its energy conversion. Most of the chemical reactions in the body are facilitated by enzymes.

Figure 2.18 Enzymes and Activation Energy Enzymes decrease the activation energy required for a given chemical reaction to occur. (a) Without an enzyme, the energy input needed for a reaction to begin is high. (b) With the help of an enzyme, less energy is needed for a reaction to begin.

2.4 Inorganic Compounds and Solutions

A Brief Introduction on Inorganic vs Organic Compounds

The Nutrition Facts label on food items breaks down the amount of fat, carbohydrates, fiber, protein, cholesterol, vitamins, and other components of the food in the package (Figure 2.19). By now you know that these components are built by atoms of various elements held together by chemical bonds. These compounds are important for the body’s structure and function. Depending on their chemical composition, the compounds are classified as either inorganic or organic. We will discuss them in detail in the last section of this chapter.

Figure 2.19 Nutrition Facts label of cereal (By BruceBlaus — Own work, CC BY-SA 4.0, https://commons.wikimedia.org/w/index.php?curid=57761221)

An inorganic compound is a substance that does not contain both carbon and hydrogen. For example, NaCl is an inorganic compound. A great number of inorganic compounds do contain hydrogen atoms, such as water (H2O) and hydrochloric acid (HCl) produced by your stomach. In contrast, only a handful of inorganic compounds contain carbon atoms. Carbon dioxide (CO2) is one of the few examples.

Inorganic compounds essential to life include water, salts, acids, and bases. Salts, also known as electrolytes, have been discussed previously in Section 2.2. In this section, we will discuss water, acids, and bases.

An organic compound, then, is a substance that contains both carbon and hydrogen. Recall that carbon and hydrogen are the second and third most abundant elements in your body (Figure 2.2: Elements of the human body). You will soon discover how these two elements combine in the foods you eat, in the compounds that make up your body structure, and in the chemicals that fuel your body.

Water: Properties of Water

We have all been told that drinking water every day is good for overall health. We also know that dehydration can cause serious health problems, such as constipation, kidney stones, problems with thermoregulation, and death. However, did you know that you can have water intoxication for overconsumption of water? As much as 70 percent of an adult’s body weight is water which 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 and life as we know it.

As a polar molecule (refer back to the polar covalent bond in Figure 2.13), a water molecule easily forms a hydrogen bond with another water molecule. This property of being attracted to another similar molecule is referred to as cohesion (Figure 2.20). When you drink a glass of water, water molecules at the surface are forming hydrogen bonds with water molecules next to them and below them, and so on. Due to this attraction, water pours out rather than scatter apart when you tilt that glass. You also notice this when rain hits the windshield of your car and remains as drops.

Surface tension is the force caused by the cohesion of molecules at the surface. That’s why light objects, such as a leaf, stay afloat and do not fall through the water surface. The polarity of water molecules is also responsible for another attractive force called adhesion. Water adheres to surfaces that have any polar or charged components and is repelled by the opposite. An example of this is when you compare water accumulation on a freshly waxed car vs. an unwaxed car. On a waxed vehicle, raindrops will repel from the waxed surface and stick together or “bead” (wax is a nonpolar substance). On a non-waxed vehicle, raindrops will disperse and spread on the polar surface of the paint and not bead.

(a)

(b)

Figure 2.20 Surface Tension of Water (a) Water molecules form a droplet with spherical shape. (b) Molecules at the surface of water experience a net attraction to other water molecules through hydrogen bonds, which hold the surface of the bulk sample together. In contrast, those in the interior experience uniform attractive forces. (CC BY-SA-NC; anonymous by request. LibreTexts.)

Water in Chemical Reactions

Water plays a key role in many biochemical reactions in the human body. Two types of chemical reactions involve the creation or the consumption of water. Dehydration synthesis is a chemical reaction where one reactant gives up a hydrogen atom and another reactant gives up a hydroxyl group (-OH) to form a new covalent bond. It is called dehydration synthesis, because the formation of this new bond results in the “synthesis” of a new product, and a water molecule is released from the reactants as a byproduct through “dehydration”  (Figure 2.21a). Dehydration synthesis reactions occur in many biochemical processes where we build large organic compounds, such as proteins, from small chemical building blocks.

Hydrolysis is, in essence, a reverse of dehydration synthesis. In this type of reaction, a water molecule disrupts a compound, breaking its covalent bonds. The water is itself split into H and OH, thus it is called hydro (water) and lysis (break). One portion of the severed compound then bonds with the hydrogen atom, and the other portion bonds with the hydroxyl group (Figure 2.21b). In your body, these reactions play an important role in breaking down large organic compounds, such as starch in rice, into smaller molecules such as glucose which can be utilized by our cells as fuel.

Figure 2.21 Dehydration Synthesis and Hydrolysis Monomers, the basic units for building larger molecules, form polymers (two or more chemically-bonded monomers). (a) In dehydration synthesis, two monomers are covalently bonded in a reaction in which one gives up a hydroxyl group and the other a hydrogen atom. A molecule of water is released as a byproduct during dehydration reactions. (b) In hydrolysis, the covalent bond between two monomers is split by the addition of a hydrogen atom to one and a hydroxyl group to the other, which requires the contribution of one molecule of water.

Retention of heat (high specific heat)

Water is important in the thermoregulation of the human body. Have you heard the saying, “a watched pot never boils”? Water with its high heat capacity can absorb and dissipate heat but does not experience a significant corresponding increase in temperature. This ability to dissipate heat is the reason why any large body of water, such as lakes or oceans, remains cooler than the land or sandy beach in warm environmental conditions such as under the hot sun. For example, if you are living in a California beach town, the temperature is much more moderate vs. living 40 miles inland which experiences much higher fluctuations in temperature extremes. This is due to water’s ability to dissipate heat.

In the body, water absorbs the heat generated by chemical reactions without a resultant increase in body temperature. Moreover, when the environmental temperature soars, the water stored in the body helps keep the body cool. This cooling effect happens as warm blood from the body’s core flows to the blood vessels just under the skin and heat is transferred to the environment via radiation. At the same time, sweat glands release warm water in sweat. As the water from sweat evaporates into the air, it carries away heat, and then the cooler blood from the periphery circulates back to the body core.

Water as a polar solvent

It is believed that life cannot exist without water because water is considered a “universal solvent”. The solvent is a liquid that can dissolve other substances. Water is certainly the most abundant solvent in the body; essentially all of the body’s chemical reactions occur among compounds dissolved in water. Because water molecules are polar, water readily dissolves ionic compounds and polar covalent compounds. Such compounds are referred to as hydrophilic, which means “water (hydro-)-loving (-philic).” As mentioned above, sugar dissolves well in water. This is because sugar molecules contain regions of hydrogen-oxygen polar bonds, making them hydrophilic. Nonpolar molecules, which do not readily dissolve in water, are called hydrophobic, or “water (hydro-)-fearing (-phobic).” For example, 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 (see Figure 2.22).

Figure 2.22 Oil and water Oil and water form two separate phases in a cup. Water and oil don’t mix well because water is a polar molecule while oil is a nonpolar molecule. Stearate (By Smokefoot – Own work, CC BY-SA 3.0) and water. 

Mixtures: Solution, Colloid, and Suspension

Examples of Mixtures

A mixture is a combination of two or more substances, each of which maintains its own chemical identity. For example, consider stirring flour and sugar together in a bowl. They do not bond to form a new compound — they are still sugar and flour, but just mixed. Another example of a mixture is the room air you breathe. Air is a gaseous mixture of molecules, including nitrogen (N2) oxygen (O2), and carbon dioxide (CO2). There are three types of liquid mixtures, all of which contain water as a key component. These are solutions, colloids, and suspensions (Figure 2.23).

Figure 2.23 Examples of Mixtures (a) A solution is a homogeneous mixture that appears clear, such as the saltwater in this aquarium. (b) In a colloid, such as milk, the particles are much larger but remain dispersed and do not settle. (c) A suspension, such as mud, is a heterogeneous mixture of suspended particles that appears cloudy and in which the particles can settle. (credit a photo: modification of work by Adam Wimsatt; credit b photo: modification of work by Melissa Wiese; credit c photo: modification of work by Peter Burgess) (https://openstax.org/books/chemistry-atoms-first-2e/pages/11-5-colloids?query=colloids%20suspension&target=%7B%22index%22%3A0%2C%22type%22%3A%22search%22%7D#fs-idm25079776)

For cells in the body to survive, they must be kept moist in a water-based liquid called a solution. A liquid solution consists of a solvent that dissolves a substance called a solute. An important characteristic of solutions is that they are homogeneous; that is, the solute molecules are distributed evenly throughout the solution. For example, if you stir a teaspoon of sugar into a glass of water, the sugar would dissolve into sugar molecules separated by water molecules. If you were to add more sugar, the ratio of sugar to water would change, but the distribution — provided you had stirred well — would still be even.

A colloid is a mixture that is somewhat like a heavy solution. The solute particles consist of tiny clumps of molecules large enough to make the liquid mixture opaque (because the particles are large enough to scatter light). Familiar examples of colloids are milk and cream. In the thyroid glands, the thyroid hormone is stored as a thick protein mixture also called a colloid.

A suspension is a liquid mixture in which a heavier substance is suspended temporarily in a liquid, but over time, settles out. This separation of particles from a suspension is called sedimentation. An example of sedimentation occurs in the blood test that establishes sedimentation rate, or sed rate. The test measures how quickly red blood cells in a test tube settle out of the watery portion of blood (known as plasma) over a set period. Rapid sedimentation of blood cells does not normally happen in the healthy body, but certain pathologies such as inflammation can cause blood cells to clump together, and these heavy clumps of blood cells settle to the bottom of the test tube more quickly than do normal blood cells.

The concentration of a given solute is the number of particles of that solute in a given space. In the bloodstream of humans, glucose concentration is usually measured in milligram (mg) per deciliter (dL) or 100 mL. In a healthy adult, the fasting glucose concentration averages about 100 mg/dL, which means 100 milligrams of glucose per deciliter of blood plasma. Another method of measuring the concentration of a solute is by its molarity — which is moles (M) of the molecules per liter (L). The mole of an element is its atomic weight, while a mole of a compound is the sum of the atomic weights of its components, called the molecular weight. An often-used example is calculating a mole of glucose, with the chemical formula of glucose: C6H12O6. Using the periodic table, the atomic weight of carbon (C) is 12.011 grams (g), and there are six carbons in glucose, for a total atomic weight of 72.066 g of C. Doing the same calculations for hydrogen (H) and oxygen (O), the molecular weight equals 180.156 g (the “gram molecular weight” of glucose). When water is added to make one liter of solution, you have one mole (1M) of glucose. This is particularly useful in chemistry because of the relationship of moles to “Avogadro’s number.” A mole of any solution has the same number of particles in it: 6.02 × 1023. Many substances in the bloodstream and other tissue of the body are measured in thousandths of a mole, or millimoles (mM). For glucose, a normal fasting concentration is 5.6 mM/L. Individuals who have a fasting glucose concentration of 126 mg/dL (7 mM/L) or higher on more than one testing are considered diabetic (Cleveland Clinic).

Inorganic Acids and Bases

You have probably heard of “acids” and “bases”, but what do they mean? Acids and bases are compounds that can dissociate in water into charged ions or electrolytes. Acids and bases can very much change the properties of the solutions in which they are dissolved.

Acids

An acid is a substance that releases hydrogen ions (H+) in a solution (Figure 2.24a). Because an atom of hydrogen has no neutrons and just one proton and one electron, a positively charged hydrogen ion (having lost its only electron) is simply a proton. This solitary proton is highly likely to be reactive and participates in chemical reactions. Strong acids are compounds that tend to dissociate completely and release all of their available H+ in solution. In the human body, cells of the stomach lining produce hydrochloric acid (HCl). This is a strong acid that releases all of its H+ in the stomach’s watery environment and aids in the breakdown of food compounds and kills ingested microbes. Weak acids do not release all available H+, and some H+ remain bonded within a compound in solution. An example of a weak acid is vinegar or acetic acid.

Figure 2.24 Acids and Bases (a) In an aqueous solution, an acid dissociates into hydrogen ions (H+) and anions. Nearly every molecule of a strong acid dissociates, producing a high concentration of H+. (b) In an aqueous solution, a base dissociates into hydroxyl ions (OH–) and cations. Nearly every molecule of a strong base dissociates, producing a high concentration of OH–.

Bases

A base is a substance that releases hydroxyl ions (OH–) in solution or takes up H+ already present in solution (see Figure 2.24b). The hydroxyl ions (hydroxide ions) or other basic substances combine with H+ in the solution to form a water molecule, thereby removing free H+ and reducing the solution’s acidity. Strong bases release most or all of their hydroxyl ions; weak bases release only some hydroxyl ions or absorb only a few H+. Food mixed with HCl from the stomach next enters the small intestine. The inner lining of the small intestine is not designed to stand up to the harsh acidic environment of the stomach and therefore needs to be buffered by the release of bicarbonate (HCO3–), a weak base that attracts H+. Bicarbonate accepts some of the H+ protons, thereby reducing the acidity of the solution.

pH and the pH scale

pH is a measurement of the concentration of hydrogen ions in solution. (Figure 2.25). pH ranges on a scale from 0 to 14, with 7 being neutral — neither acidic nor basic, with 0 being the most acidic and 14 being the most basic. For example, pure water has a pH of 7 with equal amounts of H+ and OH–.  A solution’s pH is the negative, base-10 logarithm of the hydrogen ion (H+) concentration of the solution (pH = – log [H+]). This means that the concentration of H+ at each pH value is 10 times different than the next pH. As an example, a pH 1 solution has an H+ concentration that is ten times greater than that of a pH 2 solution. That means a pH 1 solution is ten times more acidic than a pH 2 solution. Therefore, battery acid (pH=0) is way more acidic than lemon juice (pH=2). Human urine samples (pH= 6), for example, are ten times more acidic than pure water (pH=7). On the other hand, the higher the number above 7, the more basic (alkaline) the solution, or the lower the concentration of H+.

Figure 2.25. The pH Scale This figure shows a vertical arrow with the top half showing the basic scale and the bottom half showing the acidic scale. Different chemicals and their pH are also shown.

Strong acids in battery acid and strong bases in liquid drain cleaner are corrosive in the sense that they will dissolve the structure of an object. If strong acids or bases were to come into contact with your skin, deterioration or chemical burns can happen immediately. This is because these strong acids and bases hydrolyze (breakdown) the biomolecules such as proteins and lipids of our cells. As a result, caution must be exercised when handling strong acids and bases by wearing personal protective equipment such as a lab coat, safety goggles, and gloves.

Buffers

The pH of human blood normally ranges from 7.35 to 7.45, which is slightly alkaline. The body’s homeostatic mechanisms normally keep the pH of blood within this narrow range. Maintaining this pH range is critical because fluctuations from this range — either too acidic or too alkaline — can lead to life-threatening disorders. For example, we exhale CO2 through breathing and eliminate excess H+ and other acids in urine. The body also releases chemicals collectively called buffers into body fluids. A buffer is a solution of a weak acid and its conjugate base that can neutralize small amounts of acids or bases in body fluids. For example, if there is even a slight decrease below 7.35 in the pH of a bodily fluid, the buffer in the fluid — in this case, acting as a weak base — will bind to and eliminate the excess H+. In contrast, if pH rises above 7.45, the buffer will act as a weak acid and releases H+.

2.5 Organic Compounds

In addition to inorganic compounds, the food and drinks that we consume contain organic compounds consisting of groups of carbon atoms covalently bonded to hydrogen, usually oxygen, and often other elements as well. Organic compounds are found everywhere including every cell of the human body. The four types most important to human structure and function are carbohydrates, lipids, proteins, and nucleic acids. Before exploring these compounds, you need to first understand the chemistry of carbon which is the essential backbone of organic compounds.

Chemistry of Carbon and Functional Groups

Why are organic compounds so ubiquitous? It is because of their carbons. Carbon atoms are what make these compounds so diverse and complex. Recall that carbon atoms have four electrons in their valence shell and that the octet rule dictates that atoms tend to react in such a way as to complete their valence shell with eight electrons (Figure 2.4b). This means that a carbon atom can form up to four covalent bonds.

Commonly, carbon atoms form covalent bonds with other carbon atoms, often forming a long carbon chain referred to as a carbon skeleton or carbon backbone. Besides forming carbon backbone, carbon atoms tend to share electrons with a variety of other elements, one of the most common interactions is with hydrogen. Carbon and hydrogen groupings are called hydrocarbons. If you study the figures of organic compounds in the remainder of this chapter, you will see several chains of hydrocarbons in one region of the compound.

Many combinations are possible to fill carbon’s four “vacancies.” Carbon may share electrons with oxygen or nitrogen or other atoms in a particular region of an organic compound. Moreover, the atoms to which carbon atoms bond may also be part of a functional group. A functional group is a group of atoms linked by strong covalent bonds and tends to function in chemical reactions as a single unit. You can think of functional groups as tightly knit “cliques” whose members are unlikely to be parted. Five functional groups are important in human physiology; these are the hydroxyl, carboxyl, amino, methyl, and phosphate groups (Table 2.1).

Table 2.1 Functional Groups Important in Human Physiology

A) Carbohydrates

The term carbohydrate means “hydrated carbon.” Recall that the root “hydro-” indicates water. A carbohydrate is a molecule composed of carbon, hydrogen, and oxygen; in most carbohydrates, hydrogen and oxygen are found in the same two-to-one relative proportions they have in the water. The chemical formula for a “generic” molecule of carbohydrate is (CH2O)n. The “n” is used to represent a number.

Carbohydrates are referred to as saccharides, a word meaning “sugars.” A monosaccharide (mono- = “one”) is the building block of carbohydrates. These building blocks, also called monomers, can be linked by covalent bonds to form a polymer, a molecule with similar units bonded together. Disaccharides (di- = “two”) are made up of two monomers. Polysaccharides (poly- = “many”) are polymers and can consist of hundreds to thousands of monomers.

Monosaccharides

Five monosaccharides are important in the body. Three of these are hexose sugars, which means they contain six atoms of carbon. These are glucose, fructose, and galactose, shown in Figure 2.26a. The remaining monosaccharides are the two pentose sugars, each of which contains five atoms of carbon (see Figure 2.26b). When you look at the atoms of these two pentoses, can you tell the difference? One of these pentoses is present in the DNA whereas the other one is present in RNA.

Figure 2.26 Five Important monosaccharides. The second carbon of ribose has a hydroxyl group whereas the second carbon of deoxyribose is missing an oxygen atom. Ribose is present in RNA and deoxyribose is present in DNA.

Disaccharides

Disaccharides are formed via dehydration synthesis of two monosaccharides, and the covalent bond linking them is referred to as a glycosidic bond (glyco- = “sugar”). Three disaccharides (shown in Figure 2.27) are important to humans. These are sucrose, commonly referred to as table sugar; lactose, or milk sugar; and maltose, or malt sugar. As you can tell from their common names, you consume these in your diet; however, your body cannot use them directly. Instead, in the digestive tract, they are split by specific enzymes into monosaccharides via hydrolysis. For example, lactose is catalyzed by lactase, an enzyme, into galactose and glucose…

Figure 2.27 Three Important Disaccharides All three important disaccharides (sucrose, lactose and maltose) are catalyzed by specific enzymes in the human body into monosaccharides.

Polysaccharides

There are three polysaccharides important to the human body (Figure 2.28). Starches are polymers of glucose. They occur in long chains called amylose or branched chains called amylopectin, both of which are stored in plant-based foods, like rice, potatoes, and wheat, and are relatively easy to digest. Glycogen is also a polymer of glucose, but it is stored in the tissues of animals, especially in the muscles and liver. It is not considered a dietary carbohydrate because very little glycogen remains in animal tissues after slaughter; however, the body stores excess glucose as glycogen, again, in the muscles and liver. Cellulose, a polysaccharide that is the primary component of the cell wall of green plants, is the component of plant food referred to as “fiber”. Unlike most mammals, termites have the ability to digest cellulose in wood because they have specific microbes living in their guts and the microbes secrete the enzymes (cellulase) to break down cellulose. In humans, cellulose/fiber is not digestible since humans don’t have the enzymes to digest them; however, dietary fiber has many health benefits. It helps you feel full so you eat less, it promotes a healthy digestive tract, and a diet high in fiber is thought to reduce the risk of heart disease and possibly some forms of cancer.

Figure 2.28 Three Important Polysaccharides Three important polysaccharides are starches, glycogen, and fiber.

2.6 Function of Carbohydrates

The human body obtains carbohydrates from plant-based foods and also synthesizes and stores carbohydrates in the form of glycogen.

Although most body cells can break down other organic compounds for fuel, all body cells use glucose. Moreover, nerve cells (neurons) in the brain, spinal cord, and through the peripheral nervous system, as well as red blood cells, are limited to using glucose for fuel. In the breakdown of glucose for energy, molecules of adenosine triphosphate, better known as ATP, are produced. Adenosine triphosphate (ATP) is composed of a ribose sugar, an adenine base, and three phosphate groups (see Figure 2.45). ATP releases free energy when its phosphate bonds are broken, and thus supplies ready energy to the cell to do work. More ATP is produced in the presence of oxygen (O2) than in pathways that do not use oxygen. The overall reaction for the conversion of the energy in glucose to energy stored in ATP can be written:

In addition to being a critical fuel source, carbohydrates are present in very small amounts in cells’ structures. For instance, some carbohydrate molecules bind with proteins to produce glycoproteins, and others combine with lipids to produce glycolipids, both of which are found in the membrane that encloses the contents of body cells.

Examples of homeostatic imbalances or clinical connections

Because glucose is an important fuel for our cells, if glucose homeostasis is thrown off balance, it can lead to serious damage to the body in the long term. Diabetes mellitus is a disease that occurs when blood glucose is too high leading to hyperglycemia. It can lead to heart disease, nerve damage, eye problems, and kidney disease (Figure 2.29). Hypoglycemia occurs when the level of glucose in the blood drops below normal. Mild hypoglycemia may result in shaking or blurred vision, but a severe case could cause seizures or convulsions.

Figure 2.29 Symptoms of Diabetes Diabetes is a disease where blood glucose level is higher than normal physiological level. (By Mikael Häggström – See above. All used images are in public domain., Public Domain,).

B) Lipids

A lipid is one of a highly diverse group of compounds made up mostly of hydrocarbons. The few oxygen atoms they contain are often at the periphery of the molecule. Their nonpolar hydrocarbons make all lipids hydrophobic or have a fear of water (see Figure 2.30). In water, lipids do not form a true solution, but they may form an emulsion, which is the term for a combination of solutions that do not mix well.

Figure 2.30 A cup with oil and water where they form two separate phases.

Triglycerides

A triglyceride is one of the most common dietary lipid groups, and the type found most abundantly in body tissues. This compound, which is commonly referred to as fats, is formed from the synthesis of two types of molecules; glycerol and fatty acids (Figure 2.31). A glycerol backbone at the core of triglycerides consists of three carbon atoms. Three fatty acids, long chains of hydrocarbons with a carboxyl group and a methyl group at opposite ends, extend from each of the carbons of the glycerol.

Figure 2.31 Triglycerides Triglycerides are composed of glycerol attached to three fatty acids via dehydration synthesis. Notice that glycerol gives up a hydrogen atom, and the carboxyl groups on the fatty acids each give up a hydroxyl group.

Triglycerides form via dehydration synthesis between glycerol and three fatty acid chains (Figure 2.31). Glycerol gives up hydrogen atoms from its hydroxyl groups at each bond, and the carboxyl group on each fatty acid chain gives up a hydroxyl group. A total of three water molecules are thereby released.

Fatty acid chains that have no double carbon bonds anywhere along their length and therefore contain the maximum number of hydrogen atoms are called saturated fatty acids. These straight, rigid chains pack tightly together and are solid or semi-solid at room temperature (Figure 2.32a). Butter and lard are examples of saturated fatty acids, as is the fat found on a steak or in your own body. In contrast, unsaturated or polyunsaturated fatty acids have one or more double carbon bonds and are kinked at that bond (Figure 2.32b). Unsaturated fatty acids are therefore unable to pack together tightly and are liquid at room temperature. Plant oils such as olive oil typically contain unsaturated fatty acids.

(c)

(d)

Figure 2.32 Fatty Acid Shapes The level of saturation of a fatty acid affects its shape. (a) Saturated fatty acid chains are straight. (b) Unsaturated fatty acid chains contain one or more double bonds and they are kinked. (c) Butter contains saturated fatty acid. (https://commons.wikimedia.org/wiki/File:Stick-of-butter-salted.jpg#/media/File:Stick-of-butter-salted.jpg)

(d) Olive oil contains unsaturated fatty acid. (By Poyraz 72 – Own work, CC BY-SA 4.0, https://commons.wikimedia.org/w/index.php?curid=44035447).

Whereas a diet high in saturated fatty acids is suspected to increase the risk of heart disease, a diet high in unsaturated fatty acids is thought to reduce the risk. This is especially true for the omega-3 unsaturated fatty acids found in cold-water fish such as salmon (see Figure 2.33). It is named omega-3 (ω-3) because the double bond started from three atoms away from the terminal methyl group in their chemical structure.

Figure 2.33 Structure of an omega-3 fatty acid: Docosahexaenoic acid (DHA)

 The most common omega−3 fatty acids found in nature.

Finally, trans fatty acids, also called trans fats, found in some processed foods, including some stick and tub kinds of margarine, are thought to be even more harmful to the heart and blood vessels than saturated fatty acids. Trans fats are created from unsaturated fatty acids (such as corn oil) when chemically treated to produce partially hydrogenated fats (and are also produced when we fry or bake or barbeque foods using fats).

Although triglycerides get a bad reputation, they are a major fuel source for the body. When you are resting or asleep, a majority of the energy used to keep you alive is derived from triglycerides stored in your fat (adipose) tissues. Triglycerides also fuel long, slow physical activity such as hiking or gardening, and contribute a modest percentage of energy for vigorous physical activity. Dietary fat also assists the absorption and transport of the fat-soluble vitamins A, D, E, and K. Additionally, stored body fat protects and cushions the body’s bones and internal organs, and acts as insulation to retain body heat.

Fatty acids are also components of glycolipids, which are sugar (“glyco-”)-fat (“-lipids”) compounds found in the cell membrane. Lipoproteins are compounds in which the hydrophobic triglycerides are packaged in protein envelopes for transport in the blood vessels.

Phospholipids

Phospholipids are similar in structure to triglycerides. However, instead of having three fatty acids, a phospholipid is generated from a diglyceride, a glycerol with just two fatty acid chains (Figure 2.34). The third binding site on the glycerol is taken up by the phosphate group, which in turn is attached to a polar “head” region of the molecule. Recall that triglycerides are nonpolar and hydrophobic. This is true for the fatty acid portion of a phospholipid compound. However, the head of a phospholipid contains charges on the phosphate groups, as well as on the nitrogen atom. These charges make the phospholipid head hydrophilic or water-loving. Therefore, phospholipids are said to be “amphipathic” (“amphi-” = both) meaning that they have hydrophobic tails and a hydrophilic head.

Figure 2.34 Other Important Lipids (a) Phospholipids are composed of two fatty acids, glycerol, and a phosphate group. (b) Sterols are ring-shaped lipids. Shown here is cholesterol. (c) Prostaglandins are derived from unsaturated fatty acids. Prostaglandin E2 (PGE2) includes hydroxyl and carboxyl groups.

Steroids

A steroid compound (referred to as a sterol) has as its foundation a set of four hydrocarbon rings bonded to a variety of other atoms and molecules (see Figure 2.34b). The steroid that makes the most important contribution to human structure and function is cholesterol, which is synthesized by the liver in humans and animals. It is also present in most animal-based foods. As with other lipids, cholesterol’s hydrocarbons make it hydrophobic; however, it has a polar hydroxyl head that is hydrophilic. Cholesterol is an important component of bile acids, compounds that help emulsify dietary fats. In fact, the word root “chole-” refers to bile. Cholesterol is also a building block of many hormones, signaling molecules that the body releases to regulate processes at distant sites. Finally, like phospholipids, cholesterol molecules are found in the cell membrane, where their hydrophobic and hydrophilic regions help regulate the flow of substances into and out of the cell.

Diets that are high in fats and cholesterol are known to increase risks of cardiovascular disease. The reason is that eating too much saturated fat from animal-based food such as beef or pork can raise the level of bad cholesterol in the blood. A high level of bad cholesterol increases the potential risk of heart diseases and stroke. This is why the American Heart Association recommends limiting the amount of saturated fats and having a balanced diet.

Prostaglandins

Like hormones, prostaglandin is one type of signaling molecule.  Prostaglandins are derived from unsaturated fatty acids (see Figure 2.34c). One reason that the omega-3 fatty acids found in fish are beneficial to the diet is that they stimulate the production of certain prostaglandins that help regulate blood pressure and inflammation, thereby reducing the risk for heart disease. Prostaglandins also sensitize nerves to pain. One class of pain-relieving medications called nonsteroidal anti-inflammatory drugs (NSAIDs), such as aspirin, works by reducing the effects of prostaglandins.

C) Proteins

When we think of proteins, we tend to think about the delicious entree that we enjoy during a meal. Or you might associate proteins with muscle tissue. But in fact, proteins are critical components of all tissues and organs. Many proteins of the body include but are not limited to the keratin in the epidermis of skin that protects underlying tissues, the collagen (rope-like fibers) found in the dermis of the skin and bones, and the meningeal layers that cover the brain and spinal cord.

Proteins are also components of many of the body’s functional chemicals, including digestive enzymes in the digestive tract, antibodies, the neurotransmitters released by neurons, and the peptide-based hormones that regulate certain body functions such as growth hormone and insulin. While carbohydrates and lipids are composed of hydrocarbons and oxygen, all proteins are similar in that they contain nitrogen (N), carbon (C), hydrogen (H), oxygen (O), and many times sulfur (S).

Amino acids and Polypeptides

Proteins are polymers made up of nitrogen-containing monomers called amino acids. An amino acid is a molecule composed of a central carbon (alpha carbon) linked to hydrogen, an amino group (-NH2; “amine” = nitrogen-containing), a carboxyl group (-COOH), and a variable side chain (-R) (Figure 2.35). In the human body, we have 20 amino acids that are encoded directly by the DNA, and these are called standard amino acids.  From just 20 standard amino acids, they contribute to nearly all of the thousands of different proteins important for human structure and function. As an analogy, think of the thousands of words that can be constructed from our 26 letter alphabet.

Figure 2.35 Structure of an Amino Acid

What distinguishes the 20 amino acids from one another is their variable group, which is referred to as a side chain or an R-group. This group can vary in size and can be polar or nonpolar, giving each amino acid its unique characteristics.

Amino acids join via dehydration synthesis to form protein polymers. The unique bond holding amino acids together is called a peptide bond (Figure 2.36). A peptide bond is a covalent bond between two amino acids that form by dehydration synthesis. A peptide is a short chain of amino acids. Strands containing fewer than about 100 amino acids are generally referred to as polypeptides rather than proteins.

Figure 2.36 Peptide Bond Different amino acids join together to form peptides, polypeptides, or proteins via dehydration synthesis. The bonds between the amino acids are peptide bonds.

The body can synthesize most of the amino acids from components of other molecules; however, nine cannot be synthesized and have to be consumed in the diet. These are known as essential amino acids. If a particular essential amino acid is not available in sufficient quantities in the body, the synthesis of proteins containing it can slow or even cease. As we include a variety of proteins in our diets, one should understand that animal based protein contains all essential amino acids that are needed for the human body. Since some plant-based foods do not contain all essential amino acids, it is vital that those who choose to eat only plant-based proteins will need to have a balanced set of proteins in the diet.

Shape of Proteins

Just as a fork cannot be used to eat soup and a spoon cannot be used to spear meat, a protein’s shape is essential to its function. A protein’s shape is determined, most fundamentally by the sequence of amino acids of which it is made (Figure 2.37a). The sequence is called the primary structure of the protein.

Figure 2.37 The Shape of Proteins (a) The primary structure is the sequence of amino acids that make up the polypeptide chain. (b) The secondary structure, which can take the form of an alpha-helix or a beta-pleated sheet, is maintained by hydrogen bonds between amino acids in different regions of the original polypeptide strand. (c) The tertiary structure occurs as a result of further folding and bonding of the secondary structure. (d) The quaternary structure occurs as a result of interactions between two or more tertiary subunits. The example shown here is hemoglobin, a protein in red blood cells which transports oxygen to body tissues.

Most proteins are twisted or folded into more complex secondary structures that form when bonding occurs between amino acids with different properties at different regions of the polypeptide. The two most common secondary structures are a spiral called an alpha-helix and a beta-pleated sheet. If you were to take a length of string and simply twist it into a spiral, it would not hold the shape. Similarly, a strand of amino acids could not maintain a stable shape without the help of hydrogen bonds and other interactions, which create bridges between different regions of the same strand (see Figure 2.37b).

The secondary structure of proteins further folds into a compact three-dimensional shape, referred to as the protein’s tertiary structure (see Figure 2.37c). A variety of interactions give rise to protein tertiary structure, such as hydrogen bonds; ionic bonds; disulfide bridges, which are bonds between the sulfhydryl (–SH) functional groups of the amino acid side chains; hydrogen bonds; and hydrophobic interactions between nonpolar side chains. All these interactions, weak and strong, combine to determine the final three-dimensional shape of the protein and its function. Often, two or more separate polypeptides bond to form an even larger protein with a quaternary structure (see Figure 2.37d). The polypeptide subunits forming a quaternary structure can be identical or different. For instance, hemoglobin, the protein found in red blood cells, is composed of four tertiary polypeptides, two of which are called alpha chains and two of which are called beta chains.

When proteins are exposed to extreme heat, acids, bases, and other substances, proteins will lose their natural three-dimensional shape. This process is called denaturation. Because form dictates function, when proteins lose their functional shape, they can no longer perform their specified task. An everyday example of protein denaturation can be observed from the changes that take place with a raw vs. cooked egg. The egg white of chicken eggs contain proteins known as albumen. The albumen of an egg is not solid or white until it is cooked (Figure 2.38). The heat from the pan breaks the hydrogen bonds that folds the protein together. Thus, it denatures the protein, radically changing its properties from viscous and opaque to solid and white. Curdling of milk when acidic lemon juice is added, or steak changing texture from raw to well-done when you grill it, are other examples of protein denaturation.

(a)

(b)

Figure 2.38. Raw egg and Protein Denaturation (a) The raw egg contains albumen in its native liquid state. (b) Cooking the raw egg by heat will denature the structure of protein into a solid and white structure. (By miya – miya’s own file, CC BY 3.0, https://commons.wikimedia.org/w/index.php?curid=3637617) (By Scurran15 – Own work, CC BY-SA 4.0, https://commons.wikimedia.org/w/index.php?curid=40659358).

There are many examples of the importance of protein shape to its function. The long, slender shape of protein strands that make up muscle tissue is essential to their ability to contract (shorten) and relax (lengthen). Bones contain long threads of a protein called collagen that acts as scaffolding upon which bone minerals are deposited. These elongated proteins are strong and durable and typically hydrophobic. In contrast, globular proteins are globes or spheres that tend to be highly reactive and are hydrophilic. The hemoglobin proteins packed into red blood cells are an example (see Figure 2.37d). Globular proteins are abundant throughout the body, playing critical roles in most body functions. Enzymes, introduced earlier as protein catalysts, are examples of the globular proteins.

Proteins as Enzymes

If you were trying to type a paper, and every time you hit a key on your laptop there was a delay of six or seven minutes before you got a response, you would probably get a new laptop. Similarly, without enzymes to catalyze chemical reactions, the human body would be nonfunctional because it would take too long to carry out metabolism.

Enzymatic reactions — chemical reactions catalyzed by enzymes — begin when substrates bind to the enzyme. A substrate is a reactant in an enzymatic reaction. This occurs in regions of the enzyme known as active sites (Figure 2.39). Any given enzyme catalyzes just one type of chemical reaction. This characteristic, called specificity, is because a substrate with a particular shape and electrical charge can bind only to an active site corresponding to that substrate, similar to putting together two pieces of a jigsaw puzzle.

Figure 2.39 Steps in an Enzymatic Reaction According to the induced-fit model, the active site of the enzyme undergoes conformational changes upon binding with the substrate. (a) Substrates approach active sites on enzymes. (b) Substrates bind to active sites, producing an enzyme–substrate complex. (c) Changes internal to the enzyme–substrate complexes facilitate the interaction of the substrates. (d) Products are formed and released. The enzyme returns to its original form, and is ready to facilitate another enzymatic reaction.

The binding of a substrate produces an enzyme-substrate complex. Enzymes speed up chemical reactions in part because, in the enzyme-substrate complex, the substrates are oriented toward each other in an optimal position to facilitate their interaction. This promotes increased reaction speed. The enzyme then releases the product(s) and resumes its original shape. The enzyme is then free to engage in the process again and will do so as long as the substrate remains.

Other Functions of Proteins

Advertisements for protein bars, powders, and shakes all say that protein is important in building, repairing, and maintaining muscle tissue, but proteins contribute to all body tissues, from the skin to the brain cells. Also, certain proteins act as hormones, chemical messengers that help regulate body functions, For example, growth hormone is important for skeletal growth, among other roles. The basic and acidic components of amino acids enable proteins to function as buffers in maintaining acid-base balance. Proteins attract fluid, and a healthy concentration of proteins in the blood, the cells, and the spaces between cells help ensure a balance of fluids in these various “compartments.” Moreover, proteins in the cell membrane help to transport electrolytes in and out of the cell, keeping these ions in a healthy balance (See the next section on homeostatic imbalance for an example).

Like lipids, proteins can bind with carbohydrates. They can thereby produce glycoproteins or proteoglycans, both of which have many functions in the body.

The body can use proteins for energy when carbohydrate and fat intake is inadequate, and stores of glycogen and adipose tissue become depleted. However, since there is no storage site for protein except functional tissues, using protein for energy causes tissue breakdown, and results in body wasting.

Figure 2.40. The normal CFTR protein is a channel protein that helps salt (sodium chloride) move in and out of cells as seen on the right side. A mutated CFTR on the left loses the function to transport ions and thus mucus accumulates outside the cells. (https://ecampusontario.pressbooks.pub/microbio/chapter/proteins/)

Proteins play a wide range of roles in the human body. To learn about the different types of proteins, click on this link to view an interactive tour of some of these molecular machinery such as the ribosomes or ATP synthase from the Protein Data Bank (PDB).

D) Nucleic Acids

The fourth type of organic compound important to human structure and function are nucleic acids. The most well-known example of nucleic acids is DNA, deoxyribonucleic acid, which is the blueprint of a cell. All nucleic acids contain nitrogen (N) and phosphate (P) in addition to carbon, hydrogen, and oxygen.

Nucleotide structure

Nucleic acids are a polymer made up of a type of monomer called a nucleotide. A nucleotide is composed of three subunits: pentose sugar, nitrogen-containing base, and phosphate group (Figure 2.41). Pentose sugars contain five carbons, and two types of pentose sugars found in a nucleotide are either deoxyribose or ribose. There are five types of nitrogen-containing bases, also known as nitrogenous bases, found in a nucleotide. Cytosine (C), thymine (T), uracil (U) are pyrimidines. A pyrimidine is a nitrogen-containing base with a single ring structure. Adenine (A) and guanine (G) are purines, nitrogenous bases with two rings. Specific sets of nucleotides can be assembled into nucleic acids (DNA or RNA) or the energy compound adenosine triphosphate.

Figure 2.41 Nucleotides (a) The building blocks of all nucleotides are one or more phosphate groups, a pentose sugar, and a nitrogen-containing base. (b) The nitrogen-containing bases of nucleotides. (c) The two pentose sugars of DNA and RNA.

DNA and RNA

Deoxyribonucleic acid (DNA) is a nucleic acid that stores genetic information. DNA is a polymer of deoxyribonucleotides, which contains deoxyribose (so-called because it has one less atom of oxygen than ribose) plus one phosphate group and one nitrogenous base. The “choices” of bases for DNA are adenine (A), cytosine (C), guanine (G), and thymine (T). Ribonucleic acid (RNA) is a nucleic acid that helps manifest the genetic code as protein. RNA is a polymer of ribonucleotides, composed of ribose, one phosphate group, and one nitrogenous base, but the “choices” of base for RNA are adenine (A), cytosine (C), guanine (G), and uracil (U) (instead of thymine).

The individual nucleotide monomers are chain-joined at their sugar and phosphate molecules via dehydration synthesis, forming a “backbone” from which the nitrogenous bases protrude. This chain of nucleotides is referred to as a “strand” of nucleotides linked by phosphodiester bonds. In DNA, the protruding bases from two strands form hydrogen bonds which hold the two strands together. Adenine (A) always pairs up thymine (T) while cytosine (C) base pairs with guanine (G) in DNA. These specific pairings are also called a “base pair”. The two strands twist to form a shape known as a double helix (Figure 2.42). The specific order, or sequence, of nitrogenous bases within a strand of DNA form the genes that act as a molecular code instructing cells in the assembly of amino acids into proteins. Humans have almost 22,000 genes in their DNA, locked up in the 46 chromosomes inside the nucleus of each cell (except red blood cells which lose their nuclei during development). These genes carry the genetic code to build one’s body and are unique for each individual except identical twins.

Figure 2.42 DNA In the DNA double helix, two strands are bound together by hydrogen bonds between the bases of the component nucleotides. Base pairing between AT and GC pairs. (By Yikrazuul – Own work, Public Domain, https://commons.wikimedia.org/w/index.php?curid=5141568) (By Yikrazuul – Own work, Public Domain, https://commons.wikimedia.org/w/index.php?curid=5141577)

In contrast, RNA consists of a mostly single strand of sugar-phosphate backbone studded with bases (Figure 2.43). There are three types of RNA: mRNA, tRNA and rRNA. Messenger RNA (mRNA) is a copy of the DNA and it is created when DNA is transcribed with the A, U, C and G nucleotides. The coding sequence of mRNAs are to be translated into protein sequences during protein synthesis. This process occurs at the cell’s protein manufacturing plants, called the ribosomes. Transfer RNA (tRNA) is responsible for carrying the amino acids to the ribosome and ribosomal RNA (rRNA) forms parts of the ribosome structures.

Figure 2.43 Comparison of RNA and DNA RNA are mostly single-stranded while DNA are double stranded. RNA uses the nucleotides with A, U, C, and G bases. DNA uses the nucleotides with A, T, C, and G bases.

Figure 2.44 Pyrimidine Dimer UV light can cause the formation of pyrimidine dimer in DNA. (By This W3C-unspecified vector image was created with Adobe Illustrator. – DNA_UV_mutation.gif, Public Domain, https://commons.wikimedia.org/w/index.php?curid=11367690)

ATP

The nucleotide adenosine triphosphate (ATP), is composed of a ribose sugar, an adenine base, and three phosphate groups. ATP is classified as a high-energy compound because the two covalent bonds linking its three phosphates in red (Figure 2.45) store a significant amount of chemical energy. In the body, the energy released from these high-energy bonds helps fuel the body’s activities, from muscle contraction to the transport of substances in and out of cells to anabolic reactions.

Figure 2.45 Structure of Adenosine Triphosphate (ATP) The bonds in red are the high energy phosphodiester bonds. When cleaved, energy is released to help make an unfavorable reaction possible.

When a phosphate group is cleaved from ATP, the products are adenosine diphosphate (ADP), inorganic phosphate (Pi), and a subsequent release of energy. This hydrolysis reaction can be written as:

ATP + H₂O → ADP + Pi + energy

Removal of a second phosphate leaves adenosine monophosphate (AMP) and two phosphate groups. Again, these reactions also release the energy that had been stored in the phosphodiester bonds. They are reversible, too, as when a phosphate group is added to ADP by dehydration synthesis. This reaction is called phosphorylation. The same level of energy that had been released during hydrolysis must be reinvested to power this reaction to create the new bond.

Cells can also transfer a phosphate group from ATP to another organic compound. For example, when glucose first enters a cell, a phosphate group is transferred from ATP, forming glucose 6-phosphate and ADP. Once glucose is phosphorylated in this way, it can be stored as glycogen or metabolized for immediate energy.

Chapter 2 Chemistry Summary

Chapter 2 Chemistry Quiz

Key Terms

acid

compound that releases hydrogen ions (H+) in solution

activation energy

amount of energy greater than the energy contained in the reactants, which must be overcome for a reaction to proceed

adenosine triphosphate (ATP)

nucleotide containing ribose and an adenine base that is essential in energy transfer

amino acid

building block of proteins; characterized by an amino and carboxyl functional groups and a variable side-chain

anion

atom with a negative charge

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

base

compound that accepts hydrogen ions (H+) in solution

bond

electrical force linking atoms

buffer

solution containing a weak acid or a weak base that opposes wide fluctuations in the pH of body fluids

carbohydrate

class of organic compounds built from sugars, molecules containing carbon, hydrogen, and oxygen in a 1-2-1 ratio

catalyst

substance that increases the rate of a chemical reaction without itself being changed in the process

cation

atom with a positive charge

chemical energy

form of energy that is absorbed as chemical bonds form, stored as they are maintained, and released as they are broken

colloid

liquid mixture in which the solute particles consist of clumps of molecules large enough to scatter light

compound

substance composed of two or more different elements joined by chemical bonds

concentration

number of particles within a given space

covalent bond

chemical bond in which two atoms share electrons, thereby completing their valence shells

decomposition reaction

type of catabolic reaction in which one or more bonds within a larger molecule are broken, resulting in the release of smaller molecules or atoms

denaturation

change in the structure of a molecule through physical or chemical means

deoxyribonucleic acid (DNA)

deoxyribose-containing nucleotide that stores genetic information

disaccharide

pair of carbohydrate monomers bonded by dehydration synthesis via a glycosidic bond

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

enzyme

protein or RNA that catalyzes chemical reactions

exchange reaction

type of chemical reaction in which bonds are both formed and broken, resulting in the transfer of components

functional group

group of atoms linked by strong covalent bonds that tends to behave as a distinct unit in chemical reactions with other atoms

hydrogen bond

dipole-dipole bond in which a hydrogen atom covalently bonded to an electronegative atom is weakly attracted to a second electronegative atom

inorganic compound

substance that does not contain both carbon and hydrogen

ion

atom with an overall positive or negative charge

ionic bond

attraction between an anion and a cation

isotope

one of the variations of an element in which the number of neutrons differ from each other

kinetic energy

energy that matter possesses because of its motion

lipid

class of nonpolar organic compounds built from hydrocarbons and distinguished by the fact that they are not soluble in water

macromolecule

large molecule formed by covalent bonding

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

molecule

two or more atoms covalently bonded together

monosaccharide

monomer of carbohydrate; also known as a simple sugar

neutron

heavy subatomic particle having no electrical charge and found in the atom’s nucleus

nucleotide

class of organic compounds composed of one or more phosphate groups, a pentose sugar, and a base

organic compound

substance that contains both carbon and hydrogen

peptide bond

covalent bond formed by dehydration synthesis between two amino acids

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

pH

negative logarithm of the hydrogen ion (H+) concentration of a solution

phospholipid

a lipid compound in which a phosphate group is combined with a diglyceride

phosphorylation

addition of one or more phosphate groups to an organic compound

polar molecule

molecule with regions that have opposite charges resulting from uneven numbers of electrons in the nuclei of the atoms participating in the covalent bond

polysaccharide

compound consisting of more than two carbohydrate monomers bonded by dehydration synthesis via 

glycosidic bonds

potential energy

stored energy matter possesses because of the positioning or structure of its components

product

one or more substances produced by a chemical reaction

prostaglandin

lipid compound derived from fatty acid chains and important in regulating several body processes

protein

class of organic compounds that are composed of many amino acids linked together by peptide bonds

proton

heavy subatomic particle having a positive charge and found in the atom’s nucleus

purine

nitrogen-containing base with a double ring structure; adenine and guanine

pyrimidine

nitrogen-containing base with a single ring structure; cytosine, thiamine, and uracil

radioactive isotope

unstable, heavy isotope that gives off subatomic particles, or electromagnetic energy, as it decays; also called radioisotopes

reactant

one or more substances that enter into the reaction

ribonucleic acid (RNA)

ribose-containing nucleotide that helps manifest the genetic code as protein

solution

homogeneous liquid mixture in which a solute is dissolved into molecules within a solvent

steroid

(also, sterol) lipid compound composed of four hydrocarbon rings bonded to a variety of other atoms and molecules

substrate

reactant in an enzymatic reaction

suspension

liquid mixture in which particles distributed in the liquid settle out over time

synthesis reaction

type of anabolic reaction in which two or more atoms or molecules bond, resulting in the formation of a 

larger molecule

triglyceride

lipid compound composed of a glycerol molecule bonded with three fatty acid chains

valence shell

outermost electron shell of an atom

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Sources

1. OpenStax:

https://openstax.org/books/anatomy-and-physiology/pages/2-1-elements-and-atoms-the-building-blocks-of-matter

1. OpenStax (Chemistry):

https://openstax.org/books/chemistry-atoms-first-2e/pages/1-2-phases-and-classification-of-matter

https://openstax.org/books/chemistry-atoms-first-2e/pages/4-2-covalent-bonding

1. Protein Data Bank Citation: https://www.rcsb.org/pages/policies#References
2. Others: Libretexts?
3. CDC: https://www.cdc.gov/nutrition/data-statistics/plain-water-the-healthier-choice.html
4. NIH: https://www.nccih.nih.gov/health/antioxidants-in-depth
5. NIH: https://www.niddk.nih.gov/health-information/diabetes/overview/preventing-problems/low-blood-glucose-hypoglycemia
6. NIH: https://www.nhlbi.nih.gov/health-topics/cystic-fibrosis
7. (https://ecampusontario.pressbooks.pub/microbio/chapter/proteins/)
8. www.kumukahi.org
9. http://www.ulukau.org