Chapter 2 Biochemistry Basics PDF

2 Biochemistry Basics chemical discoveries are among our most promising weapons for fighting infectious and noninfectious diseases. Bacteria release abundant and assorted chemical signals that impact our own cells as well as neighboring microbes. For example, many microbes emit compounds into their...

2 Biochemistry Basics chemical discoveries are among our most promising weapons for fighting infectious and noninfectious diseases. Bacteria release abundant and assorted chemical signals that impact our own cells as well as neighboring microbes. For example, many microbes emit compounds into their surroundings that inhibit the growth of other microbes or even kill them, yet only a small number of these naturally occurring antimicrobial chemicals have been isolated and characterized. Similarly, a number of plants and The Case of the NCLEX herbs contain pharmacologically Confused Cattle Farmer HESI active ingredients, but we have a How can killer proteins TEAS poor understanding of the nature replicate themselves? of these compounds. In essence, Scan this code or visit the biochemistry is humanity’s door Mastering Microbiology Study to understanding how pathogens Area to watch the case and find out how basic biochemistry can work, the intricacies of our own explain this medical mystery. physiology, and the mechanism of action of many medical therapies. What Will We Explore? In this chapter we review the basics of atoms, molecules, and chemical reactions. We also review important macromolecules that will be discussed throughout the rest of the book. Why Is It Important? Cells carry out millions of chemical reactions every second, relying on chemistry for their very existence. Not only do we rely on chemistry for our own physiology, but it also impacts everything around us, from the tiniest bacterium to the batteries in our electronic devices. Indeed, bio- CLINICAL CASE 36 M02_NORM8290_01_SE_C02.indd 36 30/11/17 3:34 PM FROM ATOMS TO MACROMOLECULES What are atoms? Chemistry is the branch of science that studies atoms and molecules and how they interact. Biochemistry is the study of chemistry in living systems; it underpins how microbes live, how they impact us, and how we diagnose and treat the infections they cause. Ordinary matter is everything around you. It exists in five different states: solids, liquids, gases, plasmas, and Bose-Einstein condensates (discovered in 1995). Atoms are the smallest units of elements, which are pure substances that make up ordinary matter. At the center of an atom is a nucleus that contains protons and neutrons (FIG. 2.1). Protons are positively charged particles, while neutrons are noncharged or neutral particles. Around the nucleus is a cloud of negatively charged particles called electrons. Atoms can vary their number of neutrons and/or number of electrons. However, the number of protons in an atom of a given element remains constant and is a defining feature of an element. Therefore, elements are identified by their atomic number, or number of protons. Each element has a unique atomic number that is used to organize the elements into the periodic table. The Periodic Table To date, 118 elements have been identified, most of which are naturally occurring. All of the known elements are organized by atomic number, electron arrangements, and chemical properties into a table format called the periodic table of elements (FIG. 2.2). Each element is noted by its chemical symbol, an abbreviated letter notation that derives from the name of the element, which is often Greek or Latin. For example, gold is “aurum” in Latin and is noted by the symbol Au on the periodic table. Above the chemical symbol is the atomic number, which is equal to the number of protons in an atom of that given element. Underneath the chemical symbol is a number noted in smaller font––usually it is a decimal; this number is the atomic mass. Since electrons have negligible mass, the atomic mass is mainly determined by the mass of the protons and neutrons in the atom. The atomic mass is the average mass of 6.022 * 1023 atoms, or one mole, of the given element.1 After reading this section, you should be able to:   1 Define the term atom and describe its parts.   2 Determine the atomic mass, atomic number, and chemical symbol of an element using the periodic table.   3 Describe the difference between an anion and a cation and state how they are formed.   4 Discuss what isotopes are and explain how they are important in medicine.   5 Define the terms molecule, compound, and isomer.   6 Interpret and write a molecular formula.   7 Differentiate between organic and inorganic compounds and recognize selected functional groups.   8 Compare acids and bases and discuss their effects on pH.   9 Summarize what pH is and list features of the pH scale. 10 Define the term buffer and state why buffers are important in biological systems. Atomic nucleus Atomic number = 6 (6 protons, 6 neutrons) Ions and Isotopes: Variations of Atoms Two forms of atoms are ions and isotopes. All elements exist as a variety of isotopes, while only certain elements form ions. Ions In their elemental state, atoms have an equal number of positively charged protons and negatively charged electrons and are therefore neutral. However, under certain conditions some atoms form ions, which are charged atoms that have an unequal number of protons and electrons (FIG. 2.3). A positive ion, or cation, is an atom that has lost electrons and consequently has an overall positive charge. A negative ion, or anion, has a negative charge as a result of gaining electrons. Ions are chemically noted by a superscript that spec­ ifies the ion’s charge. A calcium ion is noted as Ca2+ because it has an overall charge of + 2 as a result of losing two electrons; chlorine often exists as an anion 1 A mole is technically a collection of 6.022 * 1023 of a defined entity, or Avogadro’s number of the substance being considered. If we had 6.022 * 1023 dollar bills we could say we had a mole of dollars. However, this number is so large it would be unrealistic; you’d have to spend one billion dollars per second for 19 million years to spend it all. As such, chemists and physicists use this standard unit to refer to the quantity of very small things like atoms, molecules, and subatomic particles. If there were 6.022 * 1023 molecules of water in a sample, then we would say there was one mole of water in the sample. 6 6 Carbon atom + Electron: negative charge, negligible mass Proton: positive charge, 1 atomic mass unit Neutron: neutral charge, 1 atomic mass unit FIGURE 2.1 Basic structure of an atom A generalized carbon atom is shown. The nucleus of an atom contains protons and neutrons; electrons are found in shells around the nucleus. The mass of an atom is mainly due to the number of protons and neutrons. Atoms are actually three-dimensional structures with a cloud of electrons around the atomic nucleus. From Atoms to Macromolecules M02_NORM8290_01_SE_C02.indd 37 37 30/11/17 3:34 PM Atomic number 1 1 H H Hydrogen 1.008 Hydrogen 3 4 Li Be Lithium 6.997 2 He Chemical symbol Helium 4.003 Name 1.008 5 Atomic mass Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon 9.012 10.821 12.011 14.007 15.999 18.998 20.180 14 15 16 13 Sodium Magnesium Aluminium Silicon 22.990 24.305 26.982 28.086 K Ca 20 Sc Ti 22 23 V Cr Mn Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron 39.098 40.078 44.956 47.867 50.942 51.996 54.938 55.845 21 Cobalt Nickel 58.933 58.693 Nb Mo Tc Ru Rh Pd Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium 85.468 87.621 88.906 91.224 92.906 95.941 98 101.072 102.905 76 77 75 Ba Lu Hf Ta W Re Os Cesium Barium Lutetium Hafnium Tantalum Tungsten Rhenium 132.905 137.327 174.967 178.492 180.948 183.841 186.207 88 103 87 Fr Ra Francium Radium [223] [226] Lr 104 Rf Lawrencium Rutherfordium [262] 57 La [261] 58 105 106 107 108 Dubnium Seaborgium Bohrium Hassium [262] [266] [264] [277] 61 62 109 Mt 63 Germanium Arsenic Selenium Bromine Krypton 69.723 72.631 74.922 78.972 79.907 83.798 52 53 47 48 79 80 49 In Indium Tin Antimony Tellurium Iodine Xenon 114.818 118.710 121.760 127.603 126.904 131.293 81 Platinum Gold Mercury 195.084 196.966 200.592 112 Cn [272] 65 [285] 66 Rn Thallium Lead Bismuth Polonium Astatine Radon 204.385 207.210 208.980 [209] [210] [222] 113 114 Nihonium Flerovium [286] [289] Nh 67 Tb Dy Ho Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium 138.905 140.116 140.907 144.242 145 150.36 151.964 157.253 158.925 162.500 164.930 96 97 98 86 At Gd 95 85 Po Eu 94 84 Bi Sm 93 83 Xe Pb Pm 92 82 I 54 Te Tl 111 51 Sb Hg Rg 50 Sn Nd 91 36 Gallium Pr 90 35 Zinc Ce 89 34 65.382 Au 64 39.948 63.546 112.414 [281] Argon 35.457 Copper Meitnerium Darmstadtium Roentgenium Copernicium [268] Chlorine 32.076 Kr Cadmium 110 Sulfur 30.974 Br Silver Ds Phosphorus 33 Ar Se 107.868 Iridium 18 Cl As 106.421 192.217 32 17 S Ge Palladium 78 31 P Ga Cd 190.233 Hs 30 Si Ne Zn Ag Osmium Bh 60 46 Pt Sg 59 45 Ir Db 29 Cu Zr 74 44 28 Ni Y 73 43 27 Co 39 72 42 26 Fe 38 71 41 25 Sr 56 40 24 Rb 55 10 F Al Cs 9 O Mg 37 8 N Na 19 7 C 12 11 6 B 99 Fl 68 Er 115 Mc 116 Lv Moscovium Livermorium Tennessine [289] 69 [293] Yb Erbium Thulium Ytterbium 167.259 168.934 173.045 101 [294] 118 Og Oganesson [294] 70 Tm 100 117 Ts 102 Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium [227] 232.038 231.036 238.029 [237] [244] [243] [247] [247] [251] [252] [275] [258] [259] FIGURE 2.2 The periodic table The periodic table is organized by the chemical properties and atomic number of each element. Numbers in brackets indicate the mass of that element’s most stable isotope. Critical Thinking Make a periodic table box for a fictitious element. Include a name, chemical symbol, atomic number, atomic mass, and describe how many protons, neutrons, and electrons are in your element. For simplicity, assume that your element does not have any isotopes and exists in only one form. CHEM • NOTE Hydrogen is the only element with isotopes that have names that are different from their parent atom. Deuterium is 2H and tritium is 3H. 38 and is noted as Cl - to reflect its overall charge of –1 that results when a chlorine atom gains a single electron. Isotopes All elements exist as a mixture of isotopes, which are atoms with the same number of protons but different numbers of neutrons. Isotopes are denoted by their total number of protons and neutrons. For instance, about 99 percent of carbon atoms have six protons and six neutrons, and are therefore known as carbon-12 (also written as 12C). But two other isotopes exist for carbon: Carbon-13 (13C) has seven neutrons and six protons, while carbon-14 (14C) has eight neutrons and six protons. Carbon-12 and carbon-13 are stable nonradioactive isotopes, but carbon-14 is a radioactive isotope that decays over time. Radioactive atoms release energy (radiation energy) as their unstable nucleus breaks down. It is possible for an atom to exist as an isotope or as an ion of a given element. There are also cases where an atom simultaneously exists as both an ion and an isotope. For example, 3He is an isotope of helium, He2+ is an ion of helium, and 3He2+ is both an ion and an isotope of helium. CHAPTER 2 • Biochemistry Basics M02_NORM8290_01_SE_C02.indd 38 30/11/17 3:34 PM FIGURE 2.3 Ion formation Ions form when an atom gains or loses one of more electrons. Cations form when an atom loses one or more electrons; anions form when an atom gains one or more electrons. Forming a cation Electron lost 3 3 Cation: atom now has positive charge Lithium (Li) (neutral) 3 protons, 3 electrons Li+ cation (positive) 3 protons, 2 electrons Now more protons than electrons Critical Thinking Assume two elements, X and Y, interact to transfer electrons. In the interaction, three electrons are transferred from X to Y. Based on this information, write the correct ion representation for X and Y. CHEM • NOTE Forming an anion Electron gained 17 17 Anion: atom now has negative charge Chlorine (Cl) (neutral) 17 protons, 17 electrons Cl¯ anion (negative) 17 protons, 18 electrons Now more electrons than protons Oxidation reactions remove electrons from an atom. Reduction reactions add electrons to an atom. Oxidation and reduction reactions occur together as partners and are often termed redox reactions (red = reduction; ox = oxidation). Redox reactions make ions. We will refer to redox reactions again in the metabolism chapter. Isotopes have many applications. The rate of carbon-14 decay is used to determine the age of an organic sample in a process called radiocarbon dating. A field called nuclear medicine has also become increasingly important in modern health care. This branch of health care uses radiopharmaceuticals, or drugs that contain specific isotope formulations, to diagnose and treat certain diseases. Examples of nuclear medicine techniques include positron emission tomography (PET) scans, iodine-123 imaging to assess the thyroid gland, and radionuclide therapy to treat certain cancers. What are molecules? A molecule is formed when two or more atoms bond together. An elemental hydrogen molecule contains two atoms of hydrogen (H2), while a molecule of water contains two hydrogen atoms and one oxygen atom (H2O). Sometimes the word compound is used to describe molecules that are made of more than one type of element. For example, water is a molecule, but more specifically, water is a compound. Molecular Formulas Molecules are often noted by their molecular formula (or chemical formula), which is basically their atomic recipe. Molecular formulas reveal what elements are in a molecule as well as their ratios. For example, H2 and H2O are molecular formulas. To maintain consistency, there are rules for writing out molecular formulas. If carbon is present it is typically listed first by its chemical symbol, “C,” followed by hydrogen, and then any other elements are noted by their chemical symbol in alphabetical order. If carbon is not present, then alphabetical order by chemical symbol is usually followed. For ionic compounds, the positive ion is ordinarily listed first, followed by the negative ion. For example, hydrochloric From Atoms to Macromolecules M02_NORM8290_01_SE_C02.indd 39 39 30/11/17 3:34 PM CHEM • NOTE Chemists often use R or R-group to denote the remainder of an organic molecule. This shorthand approach focuses us on the part of the molecule being discussed instead of including cumbersome chemical structures. acid is always formulaically presented as HCl, as it is made of H + cations and Cl - anions. As with all rules, there are exceptions, but understanding the methodology of molecular formulas can help you note molecular characteristics. It is possible to have different molecule structures with the same molecular formula. These are called isomers. The majority of biological molecules have at least one isomer. The formula C6H12O6 is the molecular formula for three structurally distinct sugars: glucose, fructose, and galactose. Because fructose and galactose are isomers of glucose, they can easily be converted to glucose in our body and by many microbes as well. Organic versus Inorganic Molecules Organic molecules contain carbon and hydrogen. Inorganic molecules may contain carbon, but will lack the associated hydrogen. A classic example of an inorganic molecule that contains carbon is carbon dioxide (CO2). Organic molecules are typically more complex than inorganic molecules, but that doesn’t mean they are more important to life than inorganic compounds. Both organic and inorganic molecules are necessary for life. After all, we wouldn’t last long without the inorganic molecule oxygen (O2), nor could we exist without complex organic molecules like our genetic material, DNA (deoxyribonucleic acid). Functional Groups We classify and name organic molecules based in part on the functional groups they contain. Functional groups are molecules with shared chemical properties; they often participate in chemical reactions. There are hundreds of functional groups, and the presence of certain ones can be used to predict chemical properties of a molecule. For example, alcohols—like ethanol, isopropanol, and methanol—all contain an alcohol group (a hydroxyl TABLE 2.1 Selected Biologically Important Functional Groups Functional Group Formula Alcohol O H R Carboxyl H N Found in all alcohols and added to steroids to make sterols; the suffix “ol” on a molecule name often means an alcohol group is present (examples: cholesterol, ethanol, glycerol); alcohol groups (also called hydroxyl groups) are OH groups tagged onto organic molecules. These are not to be confused with the hydroxide ion OH - , which is an inorganic ion that is not bonded to carbon and instead tends to be free in a solution. or R-OH R Amine Notes  Note: R = the remainder of a molecule; often a carbon-based addition to the molecule Important in many organic molecules including amino acids and the nitrogen bases of nucleotides or R-NH2 H Found in a variety of organic acids such as amino acids and fatty acids; it’s considered an acid because it ionizes to form R-COO - and release H + O or R-COOH C OH R Ester O R’ C O R Ether O Common in many organic molecules especially in hydrocarbon chains; added to DNA to regulate gene expression (DNA methylation) H R C H In biology esters tend to be formed by the condensation of an alcohol and an acid, by removing water (dehydration synthesis); lipids contain esters; phospholipids in bacteria and eukaryotic cell membranes have ester linkages. Common linkage in carbohydrates; found in plasma membranes of archaea or R-O-R’ R’ R Methyl or R-COO-R’ or R-CH3 H Phosphate O R O P O2 2- or R-PO4 Found in DNA, RNA, ATP, and added to lipids, carbohydrates, and proteins; phosphate is denoted as PO43 in an inorganic form or as PO42 - when bonded to an organic molecule as a functional group O2 Sulfhydryl R S H or R-SH In cysteine and methionine (amino acids); important in building disulfide bonds in organic molecules 2 40 Do not confuse hydroxyl groups, R-OH, with hydroxide ions (OH - ); they are chemically distinct. CHAPTER 2 • Biochemistry Basics M02_NORM8290_01_SE_C02.indd 40 30/11/17 3:35 PM Acid added to water H Cl + H HCl (Hydrochloric acid) H H O Cl¯ H2O (Water) Cl¯ (Chloride ion) + H+ + H O H2O (Water) H+ (Hydrogen ion) Acid contributes H+ ion to solution CHEM • NOTE Base added to water H O H Na + NaOH (Sodium hydroxide, a base) H O H Na+ + + O H O H H2O (Water) Na+ (Sodium ion) OH– (Hydroxide ion) Base contributes OH¯ ion to solution pH is calculated based on the concentration of H + ions: pH = –log10[H + ]. The higher the concentration of H + , the lower the pH. H2O (Water) FIGURE 2.4 Acids and bases Acids contribute H + ions to a solution while bases contribute OH - ions. Critical Thinking Why would combining HCl (hydrochloric acid) and NaOH (sodium hydroxide; a base) produce water (H2O) and a salt? group denoted as R-OH), and have shared chemical properties based on the chemical features of this group.2 Selected biologically important functional groups are presented in TABLE 2.1. Acids, Bases, and Salts Most cellular chemistry occurs in an aqueous solution––that’s a liquid mixture where water is the solvent (dissolving agent), and the dissolved substance is called the solute. Acids contribute hydrogen ions (H + ) to an aqueous solution.3 Bases release hydroxide ions (OH - ) in an aqueous solution (FIG. 2.4). Salts form when acids and bases react with each other; the acid contributes the anion of a salt, while the base contributes a cation. The concentration of a solution is determined by the amount of solute dissolved in a specific volume of solvent (FIG. 2.5). In clinical settings you will encounter several measures of concentration, so let’s review them briefly: Molarity is a measure of the concentration of a given solute in a liter of solvent (mol/L). Many blood chemistry values, like potassium levels, are reported in the smaller unit millimoles (mmol/L) to reflect the concentration of a given substance in a liter of blood. A patient’s blood glucose level is also measured as a concentration and is usually presented as mg/dL (milligrams per deciliter); this is an example of a weight-volume proportion.4 Intravenous solutions are typically labeled as having a particular percentage of a given solute. For example, saline solutions are labeled with a percentage: 0.9 percent normal saline is a solution that contains 9 grams of NaCl per liter of water. In reality, H + ions do not exist free in a solution. Instead, they interact with water to form H3O + , or hydronium ions. For this reason H3O + ions, instead of H + ions, are often discussed when reviewing pH. Chemists have a number of different ways to define acids and bases. In this text we apply the more simplified (although limited) Arrhenius definition of acids and bases. 3 4 A milligram (mg) is 1/1000 of a gram (g); a deciliter (dL) is 1/10 of a liter (L) or 100 milliliters (mL). + 1 mg solute = 100 mL solvent 0.01 mg/mL or 0.001% solution FIGURE 2.5 Solutes, solvents, solutions Solutes are dissolved in solvents to make solutions. A solution’s concentration can be by molarity, a percentage, or a weight-volume proportion. Critical Thinking Assume you have been instructed to give a child an oral medication with a weight-volume concentration of 0.05 mg per mL (0.05 milligrams of the drug are dissolved per milliliter of solution). The child is to receive 0.1 mg of the drug every six hours. How many milliliters of the medication would you give the child every six hours? How many total milligrams would be administered over a 24-hour period? From Atoms to Macromolecules M02_NORM8290_01_SE_C02.indd 41 41 30/11/17 3:35 PM 1 molar hydrochloric acid (HCl) 0 Lemon juice 2 1 3 Extremely acidic (more H+ ions) 4 Rainwater 5 6 Pure water 7 Sea water 8 9 Household ammonia 10 Neutral (equal OH¯ and H+ ions) 11 12 1 molar sodium hydroxide (NaOH) 13 14 pH: A Measure of Acidity The balance of H + and OH - ions is what determines the overall acidity or basicity of a solution, also called the pH (potential ­hydrogen) of a solution. We use the pH scale to describe the acidity and basicity of a solution (FIG. 2.6). Typically pH values fall between zero and 14, but it is entirely possible to have pH values lower than zero and higher than 14, even in natural environments. For example, the hot springs around the Ebeko volcano in Russia contain hydrochloric acid and sulfuric acid and have a pH as low as –1.7.5 The logarithmic (log10) nature of the pH scale means there is a tenfold difference in H + ions for every whole-number increment on the scale. When there is a mixture of OH - and H + ions, they combine to form water in what is called a neutralization reaction. Pure water has an equal concentration of H + and OH - and is chemically neutral (pH = 7). Basic, or alkaline, solutions have a higher concentration of OH - compared to H + ions and exhibit a pH greater than 7. Acidic solutions have a higher concentration of H + than OH - ions and have a pH less than 7. When an acid is added to a basic solution, every H + released by the acid neutralizes an OH - from the base and the pH will consequently decrease. Likewise, a base can be added to an acidic solution to raise pH. A classic example of the latter is taking an antacid medication to ease the discomfort of heartburn, which is acid reflux from the stomach into the esophagus. Because pH can impact a number of biochemical reactions, it is often monitored in a variety of medical and industrial applications. In microbiology, pH indicators are chemicals frequently added to growth media to monitor if microbes make acidic, neutral, or basic by-products in certain biochemical reactions. Extremely basic (more OH¯ ions) TRAINING TOMORROW’S HEALTH TEAM FIGURE 2.6 The pH scale The pH scale reflects the concentration of H + ions to OH - ions. Critical Thinking The pH scale is based on a log10 scale. Knowing this, how many times more H+ would a solution with a pH of 6 have as compared to a solution with a pH of 8? Enzymes that Make Isomers Might Also Make Good Drug Targets In many cellular processes one isomer of a molecule is converted to another by isomerization reactions. These reactions require specialized cellular enzymes called isomerases. An example is triosephosphate isomerase (TPI). This enzyme is found in glycolysis, a metabolic pathway that humans and many prokaryotes depend on to break down sugars in order to release energy. Researchers are currently investigating the use of compounds that inhibit TPI as potential drugs to combat a number of pathogens that cause parasitic infections such as leishmaniasis or shistosomiasis, or even the bacterial infection tuberculosis. Q UE STIO N 2. 1 What potential negative effect would a drug that inhibits TPI have in humans if the drug is not engineered to be highly specific for TPI in the target pathogen? Patient with shistosomiasis 5 Nordstrom, D. K., Alpers, C. N., Ptacek, C. J., & Blowes, D. W. (2000). Negative pH and extremely acidic mine waters from Iron Mountain, California. Environmental Science and Technology, 34 (2), 245–258. 42 CHAPTER 2 • Biochemistry Basics M02_NORM8290_01_SE_C02.indd 42 30/11/17 3:35 PM One common pH indicator is phenol red, which changes from red to yellow in acidic conditions. The pH of a solution can be roughly estimated with an indicator, while a pH meter can very accurately measure the concentration of H + ions in a solution if a precise pH value is required. Most microbes grow best at a pH between 6.5 and 8.5. Our own physiology is also impacted by pH; our arterial blood has a tightly regulated pH of about 7.35–7.45. If blood strays even slightly outside of this pH range, a variety of dangerous pathologies result. Acidosis (lower than normal blood pH) and alkalosis (higher than normal blood pH) are both dangerous conditions. Because pH can greatly impact physiology, organisms tend to depend on buffers, which are compounds that stabilize pH by absorbing or releasing H + ions. Our blood pH is stabilized thanks to the presence of a number of buffers, including carbonic acid (H2CO3), which releases H + ions to lower pH, and bicarbonate (HCO3 - ), which absorbs H + ions to raise pH. BUILD YOUR FOUNDATION 1. What is an atom and what are the parts of an atom? 2. Where on the periodic table are atomic mass and atomic number noted? 3. What is the difference between a cation and an anion? How are ions formed? 4. What are isotopes and why are they medically useful? Phenol red, a pH indicator in certain media, turns yellow when bacteria make acids. 5. Is O2 considered a compound, molecule, or isomer? Explain your reasoning. 6. Write the molecular formula for methane, which contains four hydrogen atoms and one carbon atom. 7. State what makes a molecule organic and name two organic and two inorganic functional groups. 8. How are acids and bases different, and what effect does each have on pH? 9. What is pH and what are features of the pH scale? 10. How do buffers stabilize pH? QUICK QUIZ Build your foundation by answering the Quick Quiz: scan this code or visit the Mastering Microbiology Study Area to quiz yourself. CHEMICAL BONDS Electrons determine what bonds can form. Chemical bonds are the “glue” or forces that bind atoms in molecules. The types of bonds present in a given molecule depend on how electrons of the bonding participants interact. Most atoms contain electrons organized in electron shells around their atomic nucleus6. Each shell has a maximum number of electrons it can hold, with those closest to the nucleus tending to hold fewer electrons than the shells farther from the nucleus. The valence shell is the outermost shell. Valence electrons are the electrons found in the valence shell; for simplicity, they can be thought of as the electrons that typically participate in chemical reactions.7 Ultimately, the type of valence electron interaction that occurs between atoms dictates what kind of chemical bond is formed. 6 Electron shells are not physical structures; rather, this term is used here to describe regions where electrons are found. Each electron shell is organized into sub-shells and orbitals. After reading this section, you should be able to: 11 Define the term valence electron and state how these electrons relate to bonding. 12 Compare and contrast ionic and covalent bonds. 13 Describe what electrolytes are and explain why they are important in biological systems. 14 Discuss what polar covalent bonding is and how it sets the stage for hydrogen bonding. 15 Explain what hydrogen bonds are. 16 Describe van der Waals interactions. 17 Define the terms hydrophobic, hydrophilic, and amphipathic and describe how they relate to micelle formation. 7 With most elements it is fine to refer to the valence electrons as exclusively occupying the valence shell. However, the transition metals (see the d-block of a standard periodic table) do not always follow the classic definition of valence electrons, since their most reactive electrons are not always exclusively in a single, outermost shell. Chemical Bonds M02_NORM8290_01_SE_C02.indd 43 43 30/11/17 3:35 PM Valence electrons are shared in covalent bonds of water CHEM • NOTE Electrostatic forces are the attraction forces that exist between positive and negative atoms or molecules. Just as a magnet is attracted to metal, so too are positive and negative atoms or molecules attracted to one another. 1 1 + 8 1 8 1 Solid ionic compound Sodium chloride (NaCl) Hydrogen atoms (H) 1 valence electron each Na+ Cl¯ Oxygen atom (O) 6 valence electrons Water molecule (H2O) Valence electrons shared FIGURE 2.7 Valence electrons Valence electrons tend to be situated in an atom’s outermost shell. They can be lost, gained, or shared to form chemical bonds between atoms. When the valence shell of an atom is full, then the electron configuration is stable, and the atom tends to be nonreactive or inert. The noble gases (including helium, neon, argon, krypton, xenon, and radon) are all examples of elements that are usually inert due to their full valence shells. When the valence shell is not full, the atom will tend to be reactive—meaning that electrons in this shell can be gained, lost, or shared with another atom to stabilize the participating atoms (FIG. 2.7). Ionic compound dissolved in water Water molecule Ionic Bonds Na+ Cl¯ FIGURE 2.8 Ionic compounds Ionic compounds always consist of ions, regardless of if they are solids or dissolved in a solution. Top: In ionic solids, ions are locked in a matrix that restricts their movement. Bottom: When ionic solids dissolve in an aqueous solution the ions are surrounded by water and are freer to move about in the solution. Critical Thinking Generation of an electrical current requires a flow of electrical charges. Most solid ionic compounds do not conduct electricity as effectively as aqueous solutions of the same ionic compounds. Explain why this is the case. (Hint: Compare ion mobility in one state versus the other.) 44 With ionic bonds, the saying “opposites attract” holds true. An ionic bond is the electrostatic force of attraction that exists between oppositely charged ions (between cations and anions). Ionic bonds form when electrons are transferred from one atom to another to make ions. Sometimes students have a misconception that the atoms of ionic compounds only exist as ions in a solution, but this is not true. The atoms of ionic compounds always exist as ions irrespective of their physical state, be it solid, liquid, or gas. In a solid state the ions are just closely packed together into a matrix framework that restricts their movement (FIG 2.8 top). Electrolytes When ionic compounds dissolve in a solution, the ions are freer to move around, and are often called electrolytes (FIG. 2.8 bottom). The most common electrolytes are acids, bases, and salts. Physiological processes rely on many electrolytes, such as sodium (Na+), calcium (Ca2+), magnesium (Mg2+), potassium (K+), chloride (Cl-), bicarbonate (HCO3-), and hydrogen phosphate (HPO42-). Electrolyte imbalances can arise from excessive sweating, diarrhea, vomiting, kidney disease, or other pathologies. Electrolytes have important roles in regulating the nervous system, our heartbeat, overall blood volume, and general water balance. Severe electrolyte imbalances can be deadly and require prompt medical attention. Covalent Bonds A covalent bond is the electrostatic force of attraction between atoms that share one or more pairs of electrons (FIG. 2.9). Carbon’s covalent bonding properties allow biological systems to make complex organic molecules. Carbon is commonly found as the core atom of organic molecules in part because it can CHAPTER 2 • Biochemistry Basics M02_NORM8290_01_SE_C02.indd 44 30/11/17 3:35 PM TRAINING TOMORROW’S HEALTH TEAM Electrolytes Save Lives Rotaviruses are a significant cause of diarrhea in children. Before the introduction of the rotavirus vaccine in 2006, almost all children in the United States suffered from this virus before their first birthday. Worldwide, especially in poorer countries where the vaccine is still not widely available, it is estimated that rotavirus kills over 400,000 children per year. Pedialyte solutions are Rotaviruses cause a watery diarrhea that leads to commonly recommended dehydration and loss of electrolytes. The virus triggers by pediatricians to treat this pathology by generating a spike in calcium ions mild dehydration that can accompany outpatient (Ca 2+ ) inside certain cells of the small intestine. The calcium ion spike leads to a cascade of tissue effects that gastrointestinal distress. ultimately impairs water and electrolyte absorption. The virus also causes intestinal cells to release chloride ions into the bowel, which leads to an even greater loss of water and electrolytes. Most deaths from rotavirus can be prevented by oral rehydration therapies, which consist of a simple water-based solution enriched with electrolytes. Most oral rehydration therapies include a specific formulation of electrolytes, such as sodium chloride (NaCl), glucose, potassium chloride (KCl), and citrate. If a solution is not effective at stabilizing the patient, or if severe dehydration is present, then intravenous rehydration therapies can be used. QU E ST I ON 2 .2 Why is plain water insufficient for treating dehydration that can develop from gastrointestinal distress? Single covalent bond One pair of shared electrons 1 1 Hydrogen molecule (H2) H H Double covalent bond Two pairs of shared electrons 8 Oxygen molecule (O2) O O Polar covalent bond Partial positive charge (+) Partial positive charge (+) 1 1 form four covalent bonds and is one of the few elements capable of catenation, which is the ability of atoms of the same element to form long chains. Polar Covalent Bonds Sometimes atoms within a molecule have a charge; for example, the atoms of ionic compounds like the salt NaCl exist as charged ions, Na + and Cl - . In such compounds the atoms are charged because electrons have been transferred. However, electrons do not have to be fully transferred as occurs in an ionic bond in order for atoms in a molecule to take on an electrical charge. Thanks to polar covalent bonds, atoms within a molecule may take on a charge due to an unequal sharing of electrons without undergoing a full transfer of electrons. If you have ever seen how three-year-olds “share” (which is to say, not very well), then you can understand the basic principle of polar covalent bonds. These are bonds where electrons are not shared equally among the bonded atoms. Some atoms, such as oxygen (O), nitrogen (N), and fluorine (F), are especially greedy when it comes to electron sharing. These highly electro­negative atoms hog 8 the electrons of the covalent bond, and consequently they take on a partial negative charge and leave the other atom in the interaction with a partially positive charge. This asymmetric charge distribution between the participants in the covalent bond is called a dipole (Figure 2.9). Molecules with polar covalent bonds, and thus dipoles, are said to be polar. The term polar just means there are two ends, or poles, of the covalent molecule––just like the Earth has north and south 8 Conversational note for non-native speakers: “hog” or “hogging” means to act in a greedy, grasping, or possessive manner versus a generous or sharing manner. 8 Unequally shared electrons create partial charges, or dipoles 8 Partial negative charge (–) Water molecule (H2O) O H H FIGURE 2.9 Covalent bonds Covalent bonds form when atoms share electrons. Atoms that share one pair of electrons have formed a single covalent bond. A double covalent bond entails sharing two pairs of electrons. Atoms like oxygen (O), nitrogen (N), and fluorine (F) form polar covalent bonds when bonded to certain atoms like hydrogen. In polar covalent bonds unequal electron sharing leads to partial charges, or dipoles, within the molecule. Critical Thinking The molecule O2, elemental oxygen, does not have a dipole despite oxygen being involved in the bond— explain why. (Hint: Consider what is fundamental to a dipole.) Chemical Bonds M02_NORM8290_01_SE_C02.indd 45 45 30/11/17 3:35 PM CHEM • NOTE Electronegativity is the tendency of an atom to attract electrons. Oxygen (O), nitrogen (N), and fluorine (F) are examples of highly electronegative elements that hog electrons when covalently bonded with less electronegative atoms. As such, O, F, and N atoms are commonly involved in polar covalent bonds. Ammonia (NH3) (+) (+) H H (+) H N Partial negative charge Hydrogen bond (–) (+) Partial positive charge (+) H H O (–) Water (H2O) FIGURE 2.10 Hydrogen bonds Hydrogen bonds are a non-covalent electrostatic attraction between molecules with dipoles. poles. In the case of a polar molecule, the end of the molecule that monopolizes electrons is partially negative and the other pole, which hardly ever gets to hold the bonding electron(s), is partially positive. As such, every polar molecule has two poles––or has “dipoles.” The presence of dipoles in polar molecules is significant because they lay the foundation for hydrogen bonds. Hydrogen Bonds: Noncovalent Interactions Hydrogen bonds do not bind atoms into molecules. Instead, hydrogen bonds are a noncovalent electrostatic attraction between two or more molecules or within a single large molecule (e.g., hydrogen bonds within a protein). The name hydrogen bond is a bit misleading, because it implies that there is a bond between hydrogen atoms; in reality the bond is an electrostatic interaction between dipoles (FIG. 2.10). What does that mean? Recall that a dipole develops when there is unequal sharing of electrons, which often occurs when hydrogen is covalently bonded to oxygen, fluorine, or nitrogen. In a polar covalent bond, hydrogen takes on a partial positive charge because it doesn’t get to interact very much with the shared electrons. In contrast, the O, F, or N atoms take on a partial negative charge because they hog the electrons in the bond. A hydrogen bond forms when the partially positive hydrogen of one polar molecule is electrostatically attracted to the partial negative charge of a nearby O, F, or N atom that is not covalently bound to the hydrogen it is attracting. As an imperfect yet hopefully memorable analogy, you can think of hydrogen as being in a relationship with an O, F, or N atom. The hydrogen is happy enough with its partner that it won’t break the bond—but it happens to have a weak attraction to another O, F, or N that is already taken (already covalently bonded to another hydrogen). Again, the attraction is not enough to make the atoms leave their respective partners, but they are attracted to each other and associate when near each other. Hydrogen bonds can exist between two or more individual molecules, in which case they are called intermolecular hydrogen bonds. Alternatively, they can occur within a single large molecule as intramolecular hydrogen bonds. The hydrogen bonding that occurs between separate H2O molecules is an example of intermolecular hydrogen bonding. The abundant hydrogen bonding between water molecules is what gives water many of its unique and biologically important properties, such as its high surface tension, its high specific heat that allows it to remain at a fairly stable temperature unless a large amount of heat is added or removed, and its solvent properties. Large macromolecules like proteins and nucleic acids often rely on intramolecular hydrogen bonds to stabilize their structures. Van der Waals Interactions Some molecules have temporary dipoles that are not a result of hydrogen atoms bonding to O, F, or N atoms. These molecules do not form hydrogen bonds; however, they still exhibit a force of attraction called van der Waals interactions (van der Waals forces). Like hydrogen bonds, van der Waals interactions do not bind atoms into molecules and instead are considered electrostatic interactions between molecules. Van der Waals interactions are weaker than hydrogen bonds and much weaker than ionic bonds, but when added up across a complex molecule they are significant stabilizers of molecular structures. Water prefers to interact with polar molecules. Perhaps you have heard the tale of Narcissus, a mythological character who was so self-impressed that he couldn’t stop staring at his own reflection upon seeing it in a pool of water. In many ways, water molecules love themselves so 46 CHAPTER 2 • Biochemistry Basics M02_NORM8290_01_SE_C02.indd 46 30/11/17 3:35 PM much that they, too, are reluctant to interact with other molecules unless the molecules in the interaction have similar properties to water—especially in terms of polarity. Water’s polar nature results from the strong dipole that is generated by the polar covalent bonds of the molecule. The resulting dipole allows water molecules to form hydrogen bonds with each other and with other polar substances. This makes water a great solvent for dissolving polar substances, but it also means water is not particularly good at dissolving nonpolar substances, like fats. Because water is so abundant in living systems, it influences the overall chemistry within biological systems and greatly impacts the macromolecular structures of biologically important macromolecules. As such, a basic appreciation of polarity is central to your training. This concept is important for understanding things like fat-soluble versus water-soluble vitamins. Vitamins like A, D, E, and K are nonpolar and thus fat soluble and can be stored in your body. In contrast, water-soluble vitamins such as the B vitamins and vitamin C are polar. Consequently, these vitamins are easily dissolved in water and not as easily stored in humans; instead, they are readily excreted in urine. The polarity of a drug molecule also impacts its absorption and distribution in the body, as well as how it is metabolized and excreted. Drugs that are administered topically are nonpolar, so they can be absorbed into the skin and cross fatty tissues to enter the body. For example, fentanyl, a potent painkiller, can be delivered by a transdermal patch because the drug readily crosses fatty tissues due to its nonpolar properties. The Case of the Confused Cattle Farmer CLINICAL CASE NCLEX HESI TEAS Practice applying what you know clinically: scan this code or visit the Mastering Microbiology Study Area to watch Part 2 and practice for future exams. Hydrophobic, Hydrophilic, and Amphipathic Molecules Substances that are readily dissolved in water are not only described as polar, but they are also said to be hydrophilic, or water loving. Substances that are not readily dissolved in water are nonpolar or hydrophobic––water fearing (FIG. 2.11). These terms essentially describe how hydrophilic and hydrophobic molecules act when associating with water. If you add oil to water, you will notice that the oil gathers into globules and eventually settles out as a layer on top of the water. This effect is observed because water is doing everything it can to minimize its interaction with the nonpolar substance in favor of maximizing its interactions with itself, or with other polar molecules that might be present. This is because the force of attraction Water molecules between the polar molecules is greater readily interact with the polar sugar than the force of attraction between the molecules and polar and nonpolar molecules. Like disdissolve the solute to solves like; polar solutes dissolve in polar make a sugar water solution. solvents, while nonpolar solutes are dissolved in nonpolar solvents. Some molecules are not simply polar or nonpolar, but instead have both hydrophobic and hydrophilic properties, and are described as being amphipathic. Amphipathic molecules are capable of forming structures called micelles. These are assemblies of amphiWater molecules pathic molecules where the hydrophobic minimize their portion of the molecule is positioned tointeraction with nonpolar substances, ward the center of the structure, while the such as oil, and do hydrophilic region of the molecule faces not dissolve them. the aqueous environment. An example of an amphipathic molecule is detergent. The hydrophobic portion of the molecule surrounds the grease on your dishes, while the hydrophilic portion FIGURE 2.11 Polar and nonpolar molecules Polar substances dissolve in water while nonpolar substances do not. Polar molecules are described as being hydrophilic because they readily interact with water. Nonpolar molecules are said to be hydrophobic because they do not interact well with water. Chemical Bonds M02_NORM8290_01_SE_C02.indd 47 47 30/11/17 3:35 PM Micelle Lipid bilayer IN WATER IN WATER Grease Hydrophilic group Hydrophilic phosphate head Hydrophobic group Amphipathic detergent molecule Hydrophobic fatty acid tails Micelle cross section FIGURE 2.12 Amphipathic molecules­ This class of molecules has both polar and nonpolar properties. Micelles (Left) and lipid bilayers (Right) are examples of how amphipathic molecules orient themselves to prevent their hydrophobic tails from contacting water. Critical Thinking Emulsifiers are often added to salad dressings to help vinegar and oil blend better. Describe how an emulsifier might work. Amphipathic phospholipid Phospholipid bilayer allows the bound grease to be readily washed away by water (FIG. 2.12 left). A cell’s plasma membrane is a bilayer made up of amphipathic lipids called phospholipids. The phosphate-containing region of the lipid is hydrophilic and faces the aqueous environments outside and inside the cell, while the hydrophilic tail regions of the lipid are sequestered in the middle of the lipid bilayer (FIG. 2.12 right). BUILD YOUR FOUNDATION 11. Why are valence electrons important to consider in chemistry? 12. What sort of bond holds sodium (Na + ) and chloride (Cl - ) ions together in the salt sodium chloride? 13. What distinguishes a covalent bond from an ionic bond? 14. What are electrolytes and why are they biologically important? 15. What are polar covalent bonds and why are they central to hydrogen bond formation? 16. How are van der Waals interactions and hydrogen bonds similar? How are they different? Build your foundation by answering the Quick Quiz: scan this code or visit the Mastering Microbiology Study Area to quiz yourself. 17. Define the terms hydrophobic, hydrophilic, and amphipathic. QUICK QUIZ 18. How would micelles be organized if they formed in a nonpolar solution as opposed to in water? CHEMICAL REACTIONS After reading this section, you should be able to: 18 Identify the reactants and products in a chemical equation. 19 Define the term catalyst. 20 Explain synthesis, decomposition, and exchange reactions. 21 Define dehydration synthesis and hydrolysis reactions. 22 Discuss what activation energy is and how it can be lowered in biochemical reactions. 23 Compare and contrast endergonic and exergonic reactions. 24 Detail what is meant by a reversible reaction and briefly describe the concept of equilibrium. 48 Chemical reactions make and break chemical bonds. Chemical reactions involve making and/or breaking chemical bonds. By doing so, they often change the substances that participate in the reaction. The ingredients of a chemical reaction are called reactants and the substances generated as a result of the reaction are called products. Every day you encounter chemical reactions—from the combustion of gasoline in a vehicle to cooking an egg on your stove. Many chemical reactions are accompanied by observable changes, such as when cut apples start to brown due to an oxidation reaction. Chemical Equations: Recipes for Reactions A chemical equation is a written representation of a chemical reaction. It is written to depict the reactants and products of the reaction as well as any specific conditions that must be met to facilitate the reaction. The reactants are listed to the left of a horizontal arrow, products are listed to the right of the arrow, and special considerations are listed above the reaction CHAPTER 2 • Biochemistry Basics M02_NORM8290_01_SE_C02.indd 48 30/11/17 3:35 PM arrow. For example, if heat needs to be added to the reaction, it is often depicted as a small triangle, a symbol called delta (Δ), above the reaction arrow. If a catalyst is required, it also is written above the reaction arrow. A catalyst is an organic or inorganic substance that increases the rate of a reaction, but it is not used up in the reaction. The physical state of the reactants and products may also be indicated as (s) for solid, (l) for liquid, (g) for gas, or (aq) for ions in an aqueous solution. Some examples of chemical equations are examined in (FIG. 2.13). Synthesis reactions Reactants A + B C6H12O6 + C6H12O6 Sucrose-6phosphate synthase C12H22O11 + H2O Decomposition reactions A B A ∆ CaCO3 (s) + B CaO (s) + CO2 (g) Decomposition by hydrolysis A + C12H22O11 + H2O lactase + B Exchange reactions involve swapping one or more components in a compound. Sometimes these are called replacement or displacement reactions. The below are examples of exchange reactions: A + BC S AC + B (a single exchange reaction) Single exchange + B C AB + CD S AD + CB (a double exchange reaction) A C B + 2 Ag (s) + Zn(NO3)2 (aq) Double exchange A B + C D BaCl2 (aq) + NaSO4 (aq) or C6H12O6 + C6H12O6 Exchange reactions 2 AgNO3 (aq) + Zn (s) Exchange Reactions Dehydration synthesis + Water A AB S A + B OH A B Water Decomposition reactions break a substance down into simpler components. In biochemical pathways it is common to add water to break the covalent bonds in complex molecules in a process called a hydrolysis reaction (FIG. 2.14 right). Hydrolysis reactions are used to break down macromolecules like proteins, polysaccharides, lipids, and nucleic acids. A standard depiction of a decomposition reaction would be: Amino acids in growing protein 2 NH3 (g) Enzyme A B Decomposition Reactions A B Dehydration synthesis A A + B S AB B 3 H2 (g) + N2 (g) Synthesis Reactions Synthesis reactions build substances by combining two or more reactants. Typically a synthesis reaction leads to the formation of more complex molecules from simpler building blocks. When cells build proteins, nucleic acids, carbohydrates, or fats, they use an elaborate series of synthesis reactions. Sometimes building a complex organic molecule requires bringing reactants together in such a way that water is released when a covalent bond is formed. These reactions are referred to as dehydration synthesis reactions (FIG. 2.14 left). Cells use dehydration synthesis reactions to build macromolecules like proteins, polysaccharides, lipids, and nucleic acids. A specific example of a dehydration synthesis reaction is the formation of a peptide bond to build proteins, which we will review later in this chapter in our discussion of proteins. Generically, a synthesis reaction can be represented by: + Product A D + C B BaSO4 (s) + 2 NaCl (aq) FIGURE 2.13 Examples of chemical reactions Hydrolysis reaction Four amino acids in a protein Unlinked amino acid H Remove water to form a covalent bond (peptide bond) H2O Add water to break a covalent bond (peptide bond) H2O FIGURE 2.14 Dehydration and hydrolysis reactions HO H Longer protein OH Protein... H ... releases an unlinked amino acid Critical Thinking The reaction 2 H2O2 S 2 H2O + O2 is a decomposition reaction where hydrogen peroxide decomposed to water and oxygen, but it is not a hydrolysis reaction; explain why. Chemical Reactions M02_NORM8290_01_SE_C02.indd 49 49 30/11/17 3:35 PM Chemical reactions consume or release energy. Activation energy Reactant ENERGY Reactions involve collisions between atoms or molecules. In order for a reaction to occur, the collisions between the reactants have to be energetic enough to lead to a change and the reactants have to be properly oriented to interact with each other. As such, there are energetic barriers to overcome in order for a chemical reaction to occur. Even if the reactants are properly oriented, they won’t react with each other if they don’t collide with sufficient energy. The minimum amount of energy needed to get a reaction started is called the activation energy (FIG. 2.15). Under physiological conditions all reactions have an activation energy, even if so miniscule as to be almost zero. Endergonic versus Exergonic Reactions Products TIME FIGURE 2.15 Activation energy This is the minimum amount of energy needed to get a reaction started. Critical Thinking Using the pictured energy diagram as a model, draw a diagram that shows an activation energy of near zero. Some reactions will ultimately release more energy than is spent to start the reaction. These exergonic reactions make products with a lower final energy than the reactants. Other reactions use more energy than is released. These endergonic reactions make products that have a higher final energy than the reactants. In biological systems the energy released by exergonic reactions is used to fuel endergonic reactions. (Later, in Chapter 8, we will discuss how proteins called enzymes help reactions occur under physiological conditions by lowering activation energy barriers. Enzymes act as catalysts in many biochemical reactions.) Reaction Reversibility Some chemical reactions, like fuel combustion, are irreversible and are written with a unidirectional arrow that points strictly from the reactants to the products. However, in a reversible reaction, the forward and reverse reactions are both possible. Initially one reaction may occur at a higher rate than the other, but eventually the forward and reverse reactions occur at the same rate and the reaction is said to be at equilibrium. Contrary to a common student misconception, equilibrium is not a static situation where the reaction just stops. Rather, there is a dynamic and continuous forming of products and the reforming of reactants at an equal rate. This does not mean that there is an equal amount of products and reactants; it just means that the total amount of products and reactants is no longer changing. A reaction at equilibrium is depicted by a set of arrows where one arrow points toward the products and the second arrow points toward the reactants, as shown here: CH3CO2H + H2O L CH3CO2- + H3O + BUILD YOUR FOUNDATION 19. Label the reactant(s), product(s), and enzyme in the following reaction and state what type of reaction it is: catalase 2H2O2 ¡ 2H2O + O2. 20. What is a catalyst and why are they useful in biological systems? 21. What class of reaction does this equation represent: Na2CO3(aq) + CaCl2(aq) L CaCO3(s) + 2NaCl(aq)? 22. Compare dehydration synthesis reactions and hydrolysis reactions. 23. In terms of energetics, why are enzymes necessary in biological systems? Build your foundation by answering the Quick Quiz: scan this code or visit the Mastering Microbiology Study Area to quiz yourself. 50 24. What are endergonic and exergonic reactions and why are they often coupled in biological systems? QUICK QUIZ 25. How can a reaction at equilibrium still have formation of products, but not exhibit a net change in the amount of products? CHAPTER 2 • Biochemistry Basics M02_NORM8290_01_SE_C02.indd 50 30/11/17 3:35 PM BIOLOGICALLY IMPORTANT MACROMOLECULES There are four main classes of biomolecules. Carbohydrates, lipids, nucleic acids, and proteins are the four main classes of biomolecules. All are essential to life (TABLE 2.2). Biomolecules often contain multiple functional groups that contribute to the chemical properties of these complex organic molecules. Because these molecules tend to be large, they are called macromolecules. Biological macromolecules are built by a series of synthesis reactions and broken down by a series of decomposition reactions. Polymerization, the process of covalently bonding together smaller foundational units (monomers), is used to build many macromolecules. Carbohydrates include simple sugars and polysaccharides. The term carbohydrate refers to organic molecules consisting of one or more sugar monomers. These polar molecules come in many different formats and have diverse structures and functions. Single sugars are foundationally built from carbon, hydrogen, and oxygen and follow a general molecular formula (CH2O)n where there is always a 2 to 1 ratio of hydrogen and oxygen and n is often 3, 5, or 6. These single sugars can be polymerized to make larger molecules like glycogen, a polymer of glucose that is mainly stored in the liver. Carbohydrate Structures Carbohydrates are also known as polysaccharides (“many sugars”) since the smallest unit of a carbohydrate is a monosaccharide, or one sugar unit (FIG. 2.16). Examples of monosaccharides include glucose, galactose, and fructose. Disaccharides consist of two monosaccharides linked together by a glycosidic bond, which is the covalent bond formed between monosaccharides to build complex sugars. Common disaccharides include lactose and sucrose (see Table 2.2). The chemical reaction that leads to the formation of a glycosidic bond is called a dehydration synthesis reaction, since water is removed from the reactants when they come together. Glycosidic bonds can be broken by After reading this section, you should be able to: 25 Identify the four main groups of biomolecules and their building blocks. 26 Describe glycosidic bonds, peptide bonds, and phosphodiester bonds. 27 Explain the structural and functional characteristics of carbohydrates, lipids, nucleic acids, and proteins. 28 Compare and contrast deoxyribonucleotides and ribonucleotides. 29 Summarize how saturation impacts lipid characteristics. 30 Describe the four levels of protein structure. 31 State what chaperone proteins do and why they are important. Monosacchar