Thomas Eisner and the Chemical Language of Nature

1. 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

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

3. 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)

4. 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 not evident at lower levels

5. 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

7. 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. • “Atoms” are names for elemental units • About 25 different chemical elements are essential to life

8. Carbon, hydrogen, oxygen, and nitrogen make up the bulk of living matter, but there are other elements necessary for life • Trace elements are important for proper function, i.e. catalysts and enzyme cofactors Table 2.2

9. Goiters are caused by iodine deficiency • Iodine intake is .15 mg compared to calcium which is 1000 mg. Discuss with your partner why? Answer: Iodine = mostly metabolism and Calcium = Bone growth and muscle contraction Figure 2.2

10. 2.3 Elements can combine to form compounds • Chemical elements combine in fixed ratios to form compounds • Different combinations determine unique properties of each compound. • Example: sodium + chlorine  sodium chloride

11. 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

12. The nucleus is surrounded by electrons • Protons and Neutrons with mass of 1, while electrons mass of 0 (1/2000 of a proton) • 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

13. 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

14. The number of neutrons may vary • Atomic mass/mass number = protons plus neutrons • Variant forms of an element are called isotopes (different number of neutrons). Same atomic number but different atomic mass. • Some isotopes are radioactive • Atoms of each element are distinguished by a specific number of protons = Atomic number (define’s element’s unique properties. Table 2.4

15. 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. Radiation to treat cancer or hyperthyroidism • Too much radiation can cause cancer: Chernobyl, Russia. Figure 2.5A Figure 2.5B

16. 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 (innermost shell holds two electrons). • Number of electrons determines chemical reactivity of an atom

17. 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

18. Two major chemical bonds used by elements to build compounds: covalent and ionic bonds For each of the bond types, two rules satisfied • The resulting compound is electrically neutral • Outer electron orbits are filled.

19. 2.7 Ionic bonds are attractions between ions of opposite charge • When atoms gain or lose electrons, charged atoms called ions are created • Loss of electron = positive charge (cation) and gain of electron = negative charge (anion) • 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)

20. Sodium and chloride ions bond to form sodium chloride, common table salt (* charges are bonding, not atoms !!!) Na+ Cl– Figure 2.7B

21. 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

22. Molecules can be represented in many ways Table 2.8

23. 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

24. 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

25. 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

26. 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

27. 2.11 Hydrogen bonds make liquid water cohesive • Cohesion is water molecules sticking together. • Adhesion is water sticking to a surface. • Due to hydrogen bonding, water molecules can move from a plant’s roots to its leaves (transpiration) • Insects can walk on water due to surface tension created by cohesive water molecules. • Jesus Lizard Figure 2.11

28. 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 – overall temperature rises slowly when heated due to hydrogen bonds breaking • As water cools, a slight drop in temperature releases a large amount of heat from the hydrogen bonds forming • Cooler at the beach then at SRHS during summer and vice-versa during winter.

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

30. 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

31. Big picture: • Why is it so important for the evolution of life to have ice less dense then water? Discuss with partner.

32. Ice is therefore less dense than liquid water, which causes it to float (a given volume of ice has fewer water molecules than an equal volume of liquid water. • If ice sank, it would seldom have a chance to thaw • Ponds, lakes, and oceans would eventually freeze solid

33. 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

34. What is pure water? • Hint: what is water made of and what is constantly going on?

36. 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: • 0-6.9 is acidic • 7.1-14 is basic • Pure water and solutions that are neither basic nor acidic are neutral, with a pH of 7

37. pH scale H+ OH– Lemon juice; gastric juice Increasingly ACIDIC (Higher concentration of H+) Grapefruit juice • The pH scale • Biological pH ranges from 1-9 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

38. So what is acid indigestion? • Have you had it? Discuss with your partner what is going on. You may have to go back to your 7th grade digestion activity!!!

39. 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 (Mucus in the stomach is alkaline!! Other mucus produced in the body may have varying pH, including cervical mucus) • Too many wings, you have a sour stomach. Wing sauce a pH of 5. What kind of buffer do you need? • Cells are kept close to pH 7 (neutral) by buffers

40. 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

41. 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

42. 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

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

44. Balancing Equations • See Chem review Worksheet.