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2. chemical foundation of life, bio 101 fall 2014

Bio Lecture ppt
The Chemistry of Life

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2. chemical foundation of life, bio 101 fall 2014

  1. 1. The Chemistry of Life
  2. 2. Atoms Make Up All Matter • Matter – Takes up space • Energy – Ability to do work
  3. 3. Atoms Make Up All Matter • Elements are fundamental types of matter – Element cannot be broken down – Bulk elements • 25 elements essential to life • Minerals • Trace elements
  4. 4. Elements in the Human Body
  5. 5. Trace Elements Trace Element: needed for survival in very small quantities Iron Iodine Fluoride
  6. 6. Trace Elements Trace Element: needed for survival in very small quantities Iron Iodine Fluoride
  7. 7. Atoms • Smallest possible “piece” of an element • Composed of – Protons – positively charged particles, atomic number – Neutrons – uncharged particle – Electron – negatively charged particle
  8. 8. Types of Subatomic Particles Particle Charge Mass Position Electron – 0 Around Nucleus Proton + 1 In Nucleus Neutron none 1 In Nucleus
  9. 9. Atomic Number and Mass Number • Mass number: the number of protons and neutrons in the nucleus • Atomic Number: the number of protons Carbon C Atomic number Element Symbol Atomic mass 6 12.0 112
  10. 10. Isotopes Isotopes: elements with the same atomic number but different mass number Isotopes of Carbon Carbon-12 Carbon-13 Carbon-14 Electrons 6 6 6 Protons 6 6 6 Neutrons 6 7 8 Mass Number (Protons + Neutrons) 12 13 14
  11. 11. Radioisotopes • Nucleus is unstable and decays (gives of energy)
  12. 12. Example Uses of Radioisotopes Use Details Isotopic labeling the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, even different nonradioactive stable isotopes can be distinguished by mass spectrometry or infrared spectroscopy. For example, in 'stable isotope labeling with amino acids in cell culture (SILAC)' stable isotopes are used to quantify proteins. If radioactive isotopes are used, they can be detected by the radiation they emit (this is called radioisotopic labeling). Radiometric dating using the known half-life of an unstable element, one can calculate the amount of time that has elapsed since a known level of isotope existed. The most widely known example is radiocarbon dating used to determine the age of carbonaceous materials. Spectroscopy Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes, both radioactive and stable. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are 1H, 2D,15N, 13C, and 31P. Mössbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as 57Fe.
  13. 13. Carbon Dating • Carbon-14: radioisotope that decays slowly – Half-life: time for half the original concentration of an isotope to decay • C-14 can be used to “age fossils”
  14. 14. Tracers • Radioisotopes can be used to identify biologically active cells (cancer cells and goiters)
  15. 15. Tracers MRI: isotopes can be used in medical imaging to view metabolically active cells in the brain
  16. 16. Radiation Therapy • The energy given off by radioisotopes is damaging to cells and can be used to treat cancers and to treat goiters.
  17. 17. Dangers of Radioactive Isotopes FUKUSHIMA, March 11th, 2011
  18. 18. Summary of Elemental Chemistry Term Definition Element a pure chemical substance consisting of a single type of atom Atom the smallest unit that defines the chemical elements and their isotopes Atomic number the number of protons found in the nucleus of an atom of that element, and therefore identical to the charge number of the nucleus Mass number the total number of protons and neutrons (together known as nucleons) in an atomic nucleus, also called atomic mass number or nucleon number Isotope variants of a particular chemical element such that while all isotopes of a given element have the same number of protons in each atom, they differ in neutron number Atomic mass the mass of an atomic particle, sub-atomic particle, or molecule; the protons and neutrons account for almost all of the mass of an atom
  19. 19. Chemical Bonds • Chemical Bonds – How elements are hooked together • Molecule – 2 or more atoms chemically joined together – Ex. O2, Cl2, H2 • Compound – Molecule composed of 2 or more DIFFERENT atoms – Ex. NaOH, H2O, NaCl, C6H12O6
  20. 20. Compound + =
  21. 21. Chemical Bonds • Its all up to the electrons! • Electrons live in orbitals – most likely location of an electron when rotating around nucleus – Each orbital has 2 electrons - more electrons, more orbitals – Orbitals are in shells – Valence shell – outermost shell, when full, shell is stable • Most atoms DO NOT have a full shell, that’s why they can bond. • Inert Elements – Have a full outer shell and cannot bond – Noble gases (Ne, He, Ar, Xe, Kr, Rn)
  22. 22. Electron Distribution Diagrams Electron “Vacancy” in energy shell 1p 6p 7p 8p Hydrogen Carbon Nitrogen Oxygen
  23. 23. Electron Distribution Diagrams
  24. 24. Types of Bonds – Covalent Bonds • Covalent Bonds – forms when 2 atoms SHARE electrons – Nonpolar Covalent Bond – Equal share of electrons – Polar Covalent Bond – Unequal share of electrons, one atom pulls electrons more than others. • Hydrogen bonds – attractions between oppositely charged particles within a single molecule, or between molecules
  25. 25. Types of Bonds – Ionic Bonds • Ionic Bonds – forms when 1 atom “takes” an electron from another – Happens when ions of opposite charge attract each other and more negative gives up electron for bond – Very strong b/c create stability in atoms
  26. 26. Ionic Bonds: Electron Transfer
  27. 27. Ionic Bonds
  28. 28. Hydrogen Bonds • Form when partial charges between two different molecules attract one another
  29. 29. Hydrogen Bonds Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Oxygen atom slightly negative (δ−) + O + + + Water molecule Hydrogen bond Figure 2.10 Hydrogen Bonds in Water. H H Hydrogen atoms slightly positive (δ+) Polar covalent bonds a. b. c. c: © The McGraw-Hill Companies, Inc./Jacques Cornell photographer
  30. 30. O H H Polar covalent bonds Hydrogen Bonds Slightly negative end
  31. 31. Water is Essential to Life • Water Regulates Temperature – Ability to resist temperature change • Body temperature • Coastal climates
  32. 32. Water is Essential to Life • Water Regulates Temperature – Evaporation • Body temperature regulation
  33. 33. Water is Essential to Life • Many Substances Dissolve in Water – Solution = solvent + solute(s) – Hydrophilic • “water-loving” – Hydrophobic • “water-fearing”
  34. 34. Water is Essential to Life • Water is Cohesive and Adhesive – Cohesion – tendency of water molecules to stick together • Surface tension – Adhesion – tendency to form hydrogen bonds with other substances • Together responsible for transport in plants
  35. 35. Water is Essential to Life • Water Expands as It Freezes – Unusual tendency – Ice less dense than liquid water • Benefits aquatic life – Formation of ice crystals deadly • Adaptations – fur in mammals
  36. 36. Figure 2.14 Ice Floats. Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. H2O molecule Ice Liquid water
  37. 37. Ice Floats Hydrogen bonds in water Hydrogen bonds in ice
  38. 38. Water is Essential to Life • Water Participates in Life’s Chemical Reactions – Chemical reaction • Reactants • Products – Reactions happen in water – Water is either a reactant or product CH4 + 2O2  CO2 + 2H20 methane + oxygen  carbon dioxide + water
  39. 39. Chemical Reactions • Chemical Reaction – 2 or more molecules “swap” atoms to make different molecules CH4 + 2O2 CO2 + 2H2O Reactants Products 6CO2 + 6H20 C6H12O6 + 6O2 Reactants Products
  40. 40. Acids and Bases • Water disassociates into H+ and OH- • Water = Neutral Solution – H+ = OH- • Acid – Substance that adds H+ to a solution – Taste sour – Found in your stomach, orange juice, tomatoes, coffee, coca-cola – HCl, H2SO4 • Base – Substance that adds OH- to a solution – Taste bitter, feel slippery, soapy – Found in detergents, soaps, cleaners – NaOH • Buffer Systems – Pairs of weak acids and bases that help resist pH changes H2O  H+ + OH-
  41. 41. pH Scale • Measures amount of H+ ions • Ranges from 0 – 14 • 0 – 6 acids • 7 neutral • 8 – 14 bases
  42. 42. Buffers • Buffer systems regulate pH in organisms – Maintaining correct pH of body fluids critical – Buffer system • Pair of weak acid and base that resist pH changes – Carbonic acid H2CO3 H+ + HCO3 - carbonic acid bicarbonate
  43. 43. Applications of Chemistry to Biology • Ocean Acidification
  44. 44. Applications of Chemistry to Biology • Ocean Acidification – the ongoing decrease in the pH of the Earth's oceans, caused by the uptake CO2 • Effects – lower metabolic rates and immune responses of ocean life – alter ocean water’s properties allowing sound to travel further, affecting prey and predators Estimated change in sea pH caused by human created CO2.
  45. 45. Applications of Chemistry to Biology
  46. 46. Applications of Chemistry to Biology Earth formation began 4.6 BYA Moon formed 4.5 BYA First solid rock 4.4 BYA First water 4.3 BYA First evidence of life 3.8 BYA While features of self-organization and self-replication are often considered the hallmark of living systems, there are many instances of abiotic molecules exhibiting such characteristics under proper conditions. Palasek showed that self-assembly of RNA molecules can occur spontaneously due to physical factors in hydrothermal vents. It is postulated that this kind of spontaneous generation could have changed simple inorganic molecules (CO2, H2O, etc.) into organic compounds.