CHAPTER 2 The Chemical Basis of Life

CHAPTER 2The Chemical Basis of Life Modules 2.1 – 2.8

Thomas Eisner and the Chemical Language of Nature • Thomas Eisner pioneered chemical ecology • the study of the chemical language of nature • He studies how insects communicate via chemical messages

Rattlebox moths release a chemical that spiders don’t like • This spider caught a rattlebox moth and then let it go

ATOMS AND MOLECULES 2.1 The emergence of biological function starts at the chemical level • Everything an organism is and does depends on chemistry • Chemistry is in turn dependent on the arrangement of atoms in molecules • In order to understand the whole, biologists study the parts (reductionism)

Molecules and ecosystems are at opposite ends of the biological hierarchy • Each level of organization in the biological hierarchy builds on the one below it • At each level, new properties emerge

D. Organ: Flight muscle of a moth Rattlebox moth • A biological hierarchy C. Cell and tissue: Muscle cell within muscle tissue Myofibril (organelle) B. Organelle: Myofibril (found only in muscle cells) Actin Myosin Atom Figure 2.1 A. Molecule: Actin

2.2 Life requires about 25 chemical elements • A chemical element is a substance that cannot be broken down to other substances by ordinary chemical means • About 25 different chemical elements are essential to life

Carbon, hydrogen, oxygen, and nitrogen make up the bulk of living matter, but there are other elements necessary for life Table 2.2

Goiters are caused by iodine deficiency Figure 2.2

2.3 Elements can combine to form compounds • Chemical elements combine in fixed ratios to form compounds • Example: sodium + chlorine  sodium chloride

2.4 Atoms consist of protons, neutrons, and electrons • The smallest particle of an element is an atom • Different elements have different types of atoms

The nucleus is surrounded by electrons • An atom is made up of protons and neutrons located in a central nucleus 2 Protons Nucleus 2 Neutrons 2 Electrons Figure 2.4A A. Helium atom

Neutrons are electrically neutral • Each atom is held together by attractions between the positively charged protons and negatively charged electrons 6 Protons Nucleus 6 Neutrons 6 Electrons B. Carbon atom Figure 2.4B

The number of neutrons may vary • Variant forms of an element are called isotopes • Some isotopes are radioactive • Atoms of each element are distinguished by a specific number of protons Table 2.4

2.5 Connection: Radioactive isotopes can help or harm us • Radioactive isotopes can be useful tracers for studying biological processes • PET scanners use radioactive isotopes to create anatomical images Figure 2.5A Figure 2.5B

2.6 Electron arrangement determines the chemical properties of an atom • Electrons are arranged in shells • The outermost shell determines the chemical properties of an atom • In most atoms, a full outer shell holds eight electrons

Atoms whose shells are not full tend to interact with other atoms and gain, lose, or share electrons Outermost electron shell (can hold 8 electrons) Electron First electron shell (can hold 2 electrons) HYDROGEN (H) Atomic number = 1 CARBON (C) Atomic number = 6 NITROGEN (N) Atomic number = 7 OXYGEN (O) Atomic number = 8 Figure 2.6

2.7 Ionic bonds are attractions between ions of opposite charge • When atoms gain or lose electrons, charged atoms called ions are created • An electrical attraction between ions with opposite charges results in an ionic bond – + Na Cl Na Cl Na Sodium atom Cl Chlorine atom Na+ Sodium ion Cl– Chloride ion Figure 2.7A Sodium chloride (NaCl)

Sodium and chloride ions bond to form sodium chloride, common table salt Na+ Cl– Figure 2.7B

2.8 Covalent bonds, the sharing of electrons, join atoms into molecules • Some atoms share outer shell electrons with other atoms, forming covalent bonds • Atoms joined together by covalent bonds form molecules

Molecules can be represented in many ways Table 2.8

THE PROPERTIES OF WATER 2.9 Water is a polar molecule • Atoms in a covalently bonded molecule may share electrons equally, creating a nonpolar molecule • If electrons are shared unequally, a polar molecule is created

This makes the oxygen end of the molecule slightly negatively charged • The hydrogen end of the molecule is slightly positively charged • Water is therefore a polar molecule • In a water molecule, oxygen exerts a stronger pull on the shared electrons than hydrogen (–) (–) O H H (+) (+) Figure 2.9

2.10 Overview: Water’s polarity leads to hydrogen bonding and other unusual properties • The charged regions on water molecules are attracted to the oppositely charged regions on nearby molecules • This attraction forms weak bonds called hydrogen bonds Hydrogen bond Figure 2.10A

as a solid • as a liquid • as a gas • Like no other common substance, water exists in nature in all three physical states: Figure 2.10B

2.11 Hydrogen bonds make liquid water cohesive • Due to hydrogen bonding, water molecules can move from a plant’s roots to its leaves • Insects can walk on water due to surface tension created by cohesive water molecules Figure 2.11

2.12 Water’s hydrogen bonds moderate temperature • It takes a lot of energy to disrupt hydrogen bonds • Therefore water is able to absorb a great deal of heat energy without a large increase in temperature • As water cools, a slight drop in temperature releases a large amount of heat

This leads to evaporative cooling • A water molecule takes a large amount of energy with it when it evaporates Figure 2.12

2.13 Ice is less dense than liquid water • Molecules in ice are farther apart than those in liquid water Hydrogen bond ICE Hydrogen bonds are stable LIQUID WATER Hydrogen bonds constantly break and re-form Figure 2.13

If ice sank, it would seldom have a chance to thaw • Ponds, lakes, and oceans would eventually freeze solid • Ice is therefore less dense than liquid water, which causes it to float

2.14 Water is a versatile solvent • Solutes whose charges or polarity allow them to stick to water molecules dissolve in water • They form aqueous solutions Na+ – – Na+ + + Cl– Cl– – – + + – Ions in solution Salt crystal Figure 2.14

2.15 The chemistry of life is sensitive to acidic and basic conditions • A compound that releases H+ ions in solution is an acid, and one that accepts H+ ions in solution is a base • Acidity is measured on the pH scale: • 0-7 is acidic • 8-14 is basic • Pure water and solutions that are neither basic nor acidic are neutral, with a pH of 7

pH scale H+ OH– Lemon juice; gastric juice Increasingly ACIDIC (Higher concentration of H+) Grapefruit juice • The pH scale Acidic solution Tomato juice Urine NEUTRAL [H+] = [OH–] PURE WATER Human blood Seawater Neutral solution Increasingly BASIC (Lower concentration of H+) Milk of magnesia Household ammonia Household bleach Oven cleaner Basic solution Figure 2.15

Buffers are substances that resist pH change • They accept H+ ions when they are in excess and donate H+ ions when they are depleted • Buffers are not foolproof • Cells are kept close to pH 7 by buffers

2.16 Connection: Acid precipitation threatens the environment • Some ecosystems are threatened by acid precipitation • Acid precipitation is formed when air pollutants from burning fossil fuels combine with water vapor in the air to form sulfuric and nitric acids Figure 2.16A

Regulations, new technology, and energy conservation may help us reduce acid precipitation • These acids can kill fish, damage buildings, and injure trees Figure 2.16B

REARRANGEMENTS OF ATOMS 2.17 Chemical reactions rearrange matter • In a chemical reaction: • reactants interact • atoms rearrange • products result 2 H2 + O2  Figure 2.17A 2 H2O

Living cells carry out thousands of chemical reactions that rearrange matter in significant ways Beta-carotene Vitamin A (2 molecules) Figure 2.17B

CHAPTER 2 The Chemical Basis of Life