1. Life’s Chemical Basis Chapter 2 Biology Concepts and Applications, Eight Edition, by Starr, Evers, Starr. Brooks/Cole, Cengage Learning 2011. Biology, Ninth Edition, by Solomon, Berg, Martin. Brooks/Cole, Cengage Learning 2011.

2. Elements elements Substances that can’t be broken down into simpler substances by ordinary chemical reactions Each element has a chemical symbol (Example: C for carbon) Four elements (oxygen, carbon, hydrogen, and nitrogen) make up more than 96% of the mass of most organisms Calcium, phosphorus, potassium, and magnesium, are present in smaller quantities Iodine and copper are trace elements

3. Periodic Table of Elements

4. Abundance of Elements

5. 2.1 Start With Atoms  Atoms Fundamental building blocks of matter Smallest unit of an element that retains that element’s chemical properties Made up of tiny subatomic particles of matter  Nucleus Positively charged protons Uncharged neutrons (except for hydrogen)  Electrons move around the nucleus Negatively charged

6. electron nucleus Fig. 2.3, p. 22

7. Atomic Number and the Periodic Table  Every element has a fixed number of protons in the atomic nucleus (atomic number) which determines an atom’s identity and defines the element  The periodic table is a chart of the elements arranged in order by atomic number and chemical behavior  Bohr models represent the electron configurations of elements as a series of concentric rings

8. The Periodic Table

9. Elements & Isotopes  Element A pure substance consisting of atoms with the same number of protons (atomic number)  Isotopes Atoms of the same element (having the same number of protons and electrons) with varying numbers of neutrons Atoms of the same element that differ in number of neutrons (atomic weight)

10. Isotopes of Carbon

11. 2.2 Putting Radioisotopes to Use  Some isotopes are unstable and tend to break down (decay) to a more stable isotope (usually a different element)  Radioisotopes are radioactive isotopes They are not stable  Emit particles (radiation) and energy as they decay spontaneously into other elements Example: 14 C decays to 14 N when a neutron decomposes to form a proton and a fast-moving electron Radioactive decay can be detected by autoradiography, on photographic film

12. Radioactive Decay  A radioisotope decays at a constant rate into the same products Example: 14 C → 14 N Radioisotopes such as 3 H (tritium), 14 C, and 32 P can replace normal molecules and are used as tracers in research  Tracer Molecule with a detectable substance attached In medicine, radioisotopes are used for both diagnosis (such as thyroid function or blood flow) and treatment (such as cancer) PET scans

13. A PET Scan

14. The ring intercepts emissions from the labeled molecules Fig. 2.5, p. 23

15. Key Concepts: ATOMS AND ELEMENTS  Atoms are fundamental units of all matter  Protons, electrons, and neutrons are their building blocks  Elements are pure substances consisting of atoms that have the same number of protons  Isotopes are atoms of the same element that have different numbers of neutrons

16. 2.3 Why Electrons Matter  Electrons occupy orbitals (volumes of space) around the nucleus  Shell model represents orbital energy levels as successively larger circles, or shells Used to view an atom’s electron structure  The energy of an electron depends on the orbital it occupies  Electrons farther from the nucleus generally have greater energy than those closer to the nucleus

17. Shell Models

18. Electron Interactions  Atoms with unpaired electrons in their outermost shell tend to interact with other atoms They donate, accept, or share electrons to eliminate vacancies vacancy no vacancy

19. Valence Electrons  The most energetic electrons (valence electrons) occupy the valence shell, represented as the outermost concentric ring in a Bohr model  Chemical behavior of an atom is determined by the number and arrangement of its valence electrons  Atoms with full valence shells are unreactive  When the valence shell is not full, an atom tends to lose, gain, or share electrons to achieve a full outer shell  Elements in the same vertical column (group) of the periodic table have similar chemical properties

20. Electrical Charge  An atom with equal numbers of protons and electrons has no net charge  Ions (positive or negative) Atoms that have gained or lost electrons  Electronegativity Measure of how strongly an atom attracts electrons from other atoms

21. Ion Formation

22. Ions  Atoms with 1, 2, or 3 valence electrons lose electrons to other atoms  become positively charged cations  Atoms with 5, 6, or 7 valence electrons gain electrons from other atoms  become negatively charged anions  Electric charges of cations and anions provide a basis for energy transformations within the cell, transmission of nerve impulses, muscle contraction, and other biological processes

23. Compounds and Molecules  Two or more atoms may combine chemically  A chemical compound consists of atoms of two or more different elements combined in a fixed ratio  Two or more atoms joined very strongly form a stable molecule  Example: H 2 0 (water) is a molecular compound

24. Representing Molecules

25. Chemical Formulas  A chemical formula is a shorthand expression that describes the chemical composition of a substance In a simplest formula (empirical formula), subscripts give the smallest ratios for atoms in a compound (e.g. NH 2 ) A molecular formula gives the actual numbers of each type of atom per molecule (e.g. N 2 H 4 ) A structural formula shows the arrangement of atoms in a molecule (e.g. water, H—O—H)

26. Key Concepts: WHY ELECTRONS MATTER  Whether one atom will bond with others depends on the number and arrangement of its electrons

27. 2.4 What Happens When Atoms Interact?  Chemical bond Attractive force that unites atoms into a molecule  Common interactions in biological molecules: Ionic bond Covalent bond Hydrogen bond

28. Ionic Bonds  Strong association between a positive ion and a negative ion (attraction of opposite charges)

29. Fig. 2-9, p. 35 11 protons17 protons and 11 electrons Sodium (Na) 17 electrons Chlorine (Cl) 10 electrons Sodium ion (Na + ) 18 electrons Chloride ion (Cl – ) Sodium chloride (NaCl) Arrangement of atoms in a crystal of salt Ionic Bonding

30. Covalent Bonds  Two atoms share a pair of electrons  Nonpolar covalent bond Atoms share electrons equally  Polar covalent bond Electrons are shared unequally One end slightly negative, other slightly positive Polar molecule has a separation of charge

31. Covalent Bonds

32. Hydrogen Bonds  Form between a hydrogen atom and an electronegative atom Each with separate polar covalent bonds  Are not chemical bonds Do not make atoms into molecules Individually weak Collectively stabilize structures of large molecules  Water molecules interact with one another extensively through hydrogen bond formation

33. Hydrogen Bonding

34. Hydrogen Bonds

35. Key Concepts: ATOMS BOND  Atoms of many elements interact by acquiring, sharing, and giving up electrons  Ionic, covalent, and hydrogen bonds are the main interactions between atoms in biological molecules

36. 2.5 Water Molecules  Water molecules are polar Form hydrogen bonds with other polar molecules Hydrophilic substances (water-loving) Hydrophobic substances (water-dreading)

37. slight negative charge on the oxygen atom The positive and negative charges balance each other; overall, the molecule carries no charge. slight positive charge on the hydrogen atoms ++ H H O Fig. 2.11, p. 28 -

38. Water: A Polar Molecule

39. Liquid Water: Hydrogen Bonds

40. Hydrogen Bonding of Water Molecules Fig. 2-13, p. 38

41. Water’s Life-Giving Properties  Polarity gives liquid water unique properties that make life possible: Resistance to temperature changes Temperature  measure of molecular motion Internal cohesion Dissolves polar and ionic substances

42. Cohesion and Adhesion cohesion Tendency of water molecules to stick to one another, due to hydrogen bonds among molecules Tendency of molecules to resist separating from one another Major mechanism of water movement in plants Evaporation: Transition of liquid to gas adhesion The ability of water to stick to other substances, particularly those with charges on their surfaces Explains how water makes things wet

43. Water’s Cohesion

44. Water’s Solvent Properties  Solvents (water) dissolve solutes (Na +, Cl - ) spheres of hydration

45. Hydration of an Ionic Compound

46. Water Temperature: From Ice to Evaporation

47. Water and Temperature Water exists in three states, which differ in degree of hydrogen bonding: gas (vapor), liquid, and ice (crystalline) Adding heat energy makes molecules move faster (increases kinetic energy) and breaks hydrogen bonds Heat  The total kinetic energy in a sample of a substance Much of the heat energy added is used to break hydrogen bonds – less energy is available  increasing temperature

48. Fig. 2-16, p. 39 212°F 100°C (a) Steam becoming water vapor (gas) 50°C (b) Water (liquid) 32°F0°C (c) Ice (solid)

49. Key Concepts: NO WATER, NO LIFE  Life originated in water and is adapted to its properties  Water has temperature-stabilizing effects, cohesion, and a capacity to act as a solvent for many other substances  These properties make life possible on Earth

50. 2.6 Acids and Bases  Concentration: The number of molecules or ions per unit volum of a solution  pH scale Indicates hydrogen ion (H + ) concentration of a solution Ranges from 0 (most acidic) to 14 (most basic or alkaline)  At pH 7 (neutral) H + and OH – concentrations are equal

51. pH of Solutions neutral solution (pH 7) Equal concentrations of hydrogen ions and hydroxide ions (concentration of each is 10 −7 mol/L) acidic solution (pH <7) Hydrogen ion concentration is higher than hydroxide ion concentration basic solution (pH >7) Hydrogen ion concentration is lower than hydroxide ion concentration

52. Ionization of Water  In pure water, a small number of water molecules dissociate into hydrogen ions (H +) and hydroxide ions (OH − ) HOH ↔ H + + OH −  The concentrations of hydrogen ions and hydroxide ions in pure water are exactly equal  Such a solution is said to be neutral – neither acidic nor basic

53. A pH Scale

54. Acids and Bases  Acids donate H + in water More H + than OH - Acid → H + + anion; ex. HCl  H + + Cl -  Bases accept H + in water More OH - than H + Base  NaOH → Na + + OH -

55. Formation of Salts When an acid and a base are mixed in water, anions from the acid and cations from the base combine to form a salt salt Example: Sodium chloride (NaCl) is a salt in which the H + of HCl has been replaced by the cation Na + HCl + NaOH → H 2 O + NaCl

56. Salts (cont.) When a salt, acid, or base is dissolved in water, its dissociated ions (electrolytes) can conduct an electric current Animals and plants contain a variety of dissolved salts (important mineral ions) essential for fluid balance and acid–base balance Homeostatic mechanisms maintain concentrations and relative amounts of various cations and anions

57. Fig. 2.14, p. 30 battery acid drain cleaner oven cleaner bleach hair remover household ammonia toothpaste hand soap milk of magnesia baking soda phosphate detergents Tums blood, tears egg white seawater pure water corn butter milk beer bread black coffee urine, tea, typical rain orange juice tomatoes, wine bananas acid rain lemon juice cola vinegar gastric fluid

58. Buffer System  A set of chemicals (a weak acid or base and its salt) that keeps the pH of a solution stable One donates ions, the other accepts them Example: bicarbonate (HCO 3 - ) OH - + H 2 CO 3 → HCO 3 - + H 2 O HCO 3 - + H + → H 2 CO 3

59. Functions of Buffer Systems  Buffers help maintain homeostasis Homeostatic mechanisms maintain appropriate pH values Example: pH of human blood is about 7.4 and must be maintained within very narrow limits buffer Substance that resists changes in pH when an acid or base is added A buffering system includes a weak acid or a weak base

60. A Buffering System  Most biological processes proceed only within a narrow pH range, usually near neutrality Acidosis Alkalosis

61. Key Concepts: HYDROGEN IONS RULE  Life is responsive to changes in the amounts of hydrogen ions and other substances dissolved in water

62. Animation: How atoms bond

63. Animation: PET scan

64. Animation: Shell models of common elements

65. Animation: Spheres of hydration

66. Animation: Structure of water

67. Animation: The pH scale

68. Animation: The shell model of electron distribution