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The Chemistry of Life

The Chemistry of Life. Chapter 3a. Chemical Benefits and Costs. Understanding of chemistry provides fertilizers, medicines, etc. Chemical pollutants damage ecosystems. Bioremediation. Use of living organisms to withdraw harmful substances from the environment. Elements.

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The Chemistry of Life

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  1. The Chemistry of Life Chapter 3a

  2. Chemical Benefits and Costs • Understanding of chemistry provides fertilizers, medicines, etc. • Chemical pollutants damage ecosystems

  3. Bioremediation Use of living organisms to withdraw harmful substances from the environment

  4. Elements • Fundamental forms of matter • Can’t be broken apart by normal means • 92 occur naturally on Earth

  5. Most Common Elements in Living Organisms Oxygen Hydrogen Carbon Nitrogen

  6. What Are Atoms? • Smallest particles that retain properties of an element • Made up of subatomic particles: • Protons (+) • Electrons (-) • Neutrons (no charge)

  7. electron proton neutron HYDROGEN HELIUM Hydrogen and Helium Atoms Fig. 2.3, p. 22

  8. Atomic Number • Number of protons • All atoms of an element have the same atomic number • Atomic number of hydrogen = 1 • Atomic number of carbon = 6

  9. Mass Number Number of protons + Number of neutrons Isotopes vary in mass number

  10. Isotopes • Atoms of an element with different numbers of neutrons (different mass numbers) • Carbon 12 has 6 protons, 6 neutrons • Carbon 14 has 6 protons, 8 neutrons

  11. Radioisotopes • Have an unstable nucleus that emits energy and particles • Radioactive decay transforms radioisotope into a different element • Decay occurs at a fixed rate

  12. Radioisotopes as Tracers • Tracer is substance with a radioisotope attached to it • Emissions from the tracer can be detected with special devices • Following movement of tracers is useful in many areas of biology

  13. Thyroid Scan • Measures health of thyroid by detecting radioactive iodine taken up by thyroid gland normal thyroid enlarged cancerous

  14. Other Uses of Radioisotopes • Drive artificial pacemakers • Radiation therapy Emissions from some radioisotopes can destroy cells. Some radioisotopes are used to kill small cancers.

  15. What Determines Whether Atoms Will Interact? The number and arrangement of their electrons

  16. Electrons • Carry a negative charge • Repel one another • Are attracted to protons in the nucleus • Move in orbitals - volumes of space that surround the nucleus y Z X When all p orbitals are full

  17. Electron Orbitals • Orbitals can hold up to two electrons (each of opposite spin) • Atoms differ in the number of occupied orbitals • Orbitals closest to nucleus are lower energy and are filled first

  18. Shell Model • First shell • Lowest energy • Holds 1 orbital with up to 2 electrons • Second shell • 4 orbitals hold up to 8 electrons CALCIUM 20p+ , 20e-

  19. Electron Vacancies • Unfilled shells make atoms likely to react • Hydrogen, carbon, oxygen, and nitrogen all have vacancies in their outer shells CARBON 6p+ , 6e- NITROGEN 7p+ , 7e- HYDROGEN 1p+ , 1e-

  20. Chemical Bonds, Molecules, & Compounds • Bond is union between electron structures of atoms • Atoms bond to form molecules • Molecules may contain atoms of only one element - O2 • Molecules of compounds contain more than one element - H2O

  21. Chemical Bookkeeping • Use symbols for elements when writing formulas • Formula for glucose is C6H12O6 • 6 carbons • 12 hydrogens • 6 oxygens

  22. Chemical Bookkeeping • Chemical equation shows reaction Reactants ---> Products • Equation for photosynthesis: 6CO2 + 6H2O ---> + C6H12O5 + 6H2O

  23. Important Bonds in Biological Molecules Ionic Bonds Covalent Bonds Hydrogen Bonds

  24. Ion Formation • Atom has equal number of electrons and protons - no net charge • Atom loses electron(s), becomes positively charged ion • Atom gains electron(s), becomes negatively charged ion

  25. Ionic Bonding • One atom loses electrons, becomes positively charged ion • Another atom gains these electrons, becomes negatively charged ion • Charge difference attracts the two ions to each other

  26. Formation of NaCl • Sodium atom (Na) • Outer shell has one electron • Chlorine atom (Cl) • Outer shell has seven electrons • Na transfers electron to Cl forming Na+and Cl- • Ions remain together as NaCl

  27. Formation of NaCl 7mm electron transfer SODIUM ATOM 11 p+ 11 e- CHLORINE ATOM 17 p+ 17 e- CHLORINE ION 17 p+ 18 e- SODIUM ION 11 p+ 10 e- Fig. 2.10a, p. 26

  28. Covalent Bonding Atoms share a pair or pairs of electrons to fill outermost shell • Single covalent bond • Double covalent bond • Triple covalent bond

  29. Nonpolar Covalent Bonds • Atoms share electrons equally • Nuclei of atoms have same number of protons • Example: Hydrogen gas (H-H)

  30. Polar Covalent Bonds • Number of protons in nuclei of participating atoms is NOT equal • Electrons spend more time near nucleus with most protons • Water - Electrons more attracted to O nucleus than to H nuclei

  31. Hydrogen Bonding • Molecule held together by polar covalent bonds has no NET charge • However, atoms of the molecule carry different partial charges • Atom in one polar covalent molecule can be attracted to oppositely charged atom in another such molecule

  32. Examples of Hydrogen Bonds one large molecule another large molecule a large molecule twisted back on itself Fig. 2.12, p. 27

  33. Properties of Water Polarity Temperature-Stabilizing Cohesive Solvent

  34. Water Is a Polar Covalent Molecule • Molecule has no net charge • Oxygen end has a slight negative charge • Hydrogen end has a slight positive charge O H H

  35. Liquid Water + H + H + _ O H O + _ + H +

  36. Hydrophilic & HydrophobicSubstances • Hydrophilic substances • Polar • Hydrogen bond with water • Example: sugar • Hydrophobic substances • Nonpolar • Repelled by water • Example: oil

  37. Temperature-Stabilizing Effects • Liquid water can absorb much heat before its temperature rises • Why? • Much of the added energy disrupts hydrogen bonding rather than increasing the movement of molecules

  38. Evaporation of Water • Large energy input can cause individual molecules of water to break free into air • As fast (hot) molecules break free, they carry away some energy leaving behind the slower (lower temperature) ones • Evaporative water loss is used by mammals to lower body temperature

  39. Why Ice Floats • In ice, hydrogen bonds lock molecules in a lattice • Water molecules in lattice are spaced farther apart then those in liquid water • Ice is less dense than water

  40. Water Cohesion • Hydrogen bonding holds molecules in liquid water together • Creates surface tension • Allows water to move as continuous column upward through stems of plants

  41. Water Is a Good Solvent • Ions and polar molecules dissolve easily in water • When solute dissolves, water molecules cluster around its ions or molecules and keep them separated

  42. – – + + + + – + – + Na+ – + + + + – + Cl– – – + – + + + – + + – + Spheres of Hydration Fig. 2.16, p. 29

  43. Concentration • It is a measure of the number of dissolved solute particles per unit volume of solvent; it is often expressed as moles per liter (1 liter ~ 1quart). • A mole of particles represents 6.023 x 1023 particles (602,300,000,000,000,000,000,000) . • Atomic weight is the weight in grams of 1 mole of atoms

  44. The concentration of any volume of water is a constant One liter (L)of pure water has a mass of 997.05 g at 25oC. The M.W. of water is 2 x 1.008 (2H atoms) + 15.999 (1O atom) = 18.015 g per mole. Hence, 1L of water has 997.05 g/18.115 g/mole = 55.346 moles Since mass and volume change in exactly the same proportion, the molarity does not change when the volume changes.

  45. Dissociation of Water I • At room temperature (25o C) H2O molecules have a very slight tendency to break up into H+ and OH- ions (forward reaction). As these build up in solution they tend to reform water (reverse reaction) . At equilibrium the forward and reverse reactions occur at equal rates or speeds, k1 and k2:

  46. Dissociation of Water II k1 = c1 * [H2O]2 H2O + H2O → OH- + H3O+ k1 / c1 = [H2O]2 k2 = c2 * [OH-] * [H3O+] OH- + H3O+ → H2O + H2O k2 / c2 = [OH-] * [H3O+]

  47. Dissociation of Water II • The equilibrium constant for dissociation, Kd is the ratio between between backward (k2) and forward (k1) rates of reaction. The forward rate (k1) varies directly with the square of the concentration of water or [H2O]2. The reverse rate (k2) varies directly with the product of H3O+ and OH- concentrations or [H3O+ ] * [OH- ]

  48. Dissociation of Water III • At 25o C at equilibrium, when k1 = k2, it is found that 55.4 moles/liter(L) of H2O react with 55.4 moles/L of H2O to form 1.0 x 10-7 mole/L of H3O+ and 1.0 x 10-7 mole/L of OH-Thus, Kd = • k2*c1= _[H3O+] * [0H-] =[1.0 x 10-7] * [1.0 x 10-7]= k1 c2 [H2O]2 [55.4]2 = 3.26 x 10-18 = c1 (since k2 = 1) c2 k1

  49. Solubility Product of Water Since [H2O] is a constant and it is insignificantly changed by water’s dissociation, we multiply both sides of the previous equation by [55. 4]2 which equals 3063. Thus, [H3O+] * [0H-] = 3.26 x 10-18 * 3063 = 1 x 10-14 This constant , 1 x 10-14 , is the Ionization Constant of Water and is also called The Solubility Product of Water

  50. Hydrogen Ions: H+ • Unbound protons • Have important biological effects • Form when water ionizes

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