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Biochemistry Part A

Biochemistry Part A. Biochemistry- the chemistry of living things. Matter. Anything that has mass and occupies space States of matter: Solid—definite shape and volume Liquid—definite volume, changeable shape Gas—changeable shape and volume. Organic/Inorganic.

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Biochemistry Part A

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  1. Biochemistry Part A • Biochemistry- the chemistry of living things

  2. Matter • Anything that has mass and occupies space • States of matter: • Solid—definite shape and volume • Liquid—definite volume, changeable shape • Gas—changeable shape and volume

  3. Organic/Inorganic • Inorganic matter- mostly non living, but essential to living organism • *in general does not contain “C”- Carbon • Exceptions: CO, CO2 • Abundant, and represent raw materials needed to build life • Organic matter- Is living, was living, came from a living thing • *in general contains “C”- carbon

  4. Composition of Matter • Elements • Cannot be broken down by ordinary chemical means • Each has unique properties: • Physical properties • Are detectable with our senses, or are measurable • Chemical properties • How atoms interact (bond) with one another

  5. Composition of Matter • Atoms • Unique building blocks for each element • Atomic symbol: one- or two-letter chemical shorthand for each element

  6. Major Elements of the Human Body • Oxygen (O) • Carbon (C) • Hydrogen (H) • Nitrogen (N) About 96% of body mass

  7. Lesser Elements of the Human Body • About 3.9% of body mass: • Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)

  8. Trace Elements of the Human Body • < 0.01% of body mass: • Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)

  9. Atomic Structure • Determined by numbers of subatomic particles • Nucleus consists of neutrons and protons

  10. Atomic Structure • Neutrons • No charge • Mass = 1 atomic mass unit (amu) • Protons • Positive charge • Mass = 1 amu

  11. Atomic Structure • Electrons • Orbit nucleus • Equal in number to protons in atom • Negative charge • 1/2000 the mass of a proton (0 amu)

  12. Models of the Atom • Orbital model: current model used by chemists • Depicts probable regions of greatest electron density (an electron cloud) • Useful for predicting chemical behavior of atoms

  13. Models of the Atom • Planetary model—oversimplified, outdated model • Incorrectly depicts fixed circular electron paths • Useful for illustrations (as in the text)

  14. Nucleus Nucleus Helium atom Helium atom 2 protons (p+) 2 neutrons (n0) 2 electrons (e–) 2 protons (p+) 2 neutrons (n0) 2 electrons (e–) (a) Planetary model (b) Orbital model Proton Neutron Electron Electron cloud Figure 2.1

  15. Identifying Elements • Atoms of different elements contain different numbers of subatomic particles • Compare hydrogen, helium and lithium (next slide)

  16. Proton Neutron Electron Hydrogen (H) (1p+; 0n0; 1e–) Helium (He) (2p+; 2n0; 2e–) Lithium (Li) (3p+; 4n0; 3e–) Figure 2.2

  17. Identifying Elements • Atomic number = number of protons in nucleus

  18. Identifying Elements • Mass number = mass of the protons and neutrons • Mass numbers of atoms of an element are not all identical • Isotopes are structural variations of elements that differ in the number of neutrons they contain

  19. Identifying Elements • Atomic weight = average of mass numbers of all isotopes

  20. Proton Neutron Electron Hydrogen (1H) (1p+; 0n0; 1e–) Deuterium (2H) (1p+; 1n0; 1e–) Tritium (3H) (1p+; 2n0; 1e–) Figure 2.3

  21. Radioisotopes • Spontaneous decay (radioactivity) • Similar chemistry to stable isotopes • Can be detected with scanners

  22. Radioisotopes • Valuable tools for biological research and medicine • Cause damage to living tissue: • Useful against localized cancers • Radon from uranium decay causes lung cancer • Other Values of Radatiosotopes…

  23. Molecules and Compounds • Most atoms combine chemically with other atoms to form molecules and compounds • Molecule—two or more atoms bonded together (e.g., H2 or C6H12O6) • Compound—two or more different kinds of atoms bonded together (e.g., C6H12O6)

  24. Chemically Inert Elements • Stable and unreactive • Outermost energy level fully occupied or contains eight electrons

  25. (a) Chemically inert elements Outermost energy level (valence shell) complete 8e 2e 2e Helium (He) (2p+; 2n0; 2e–) Neon (Ne) (10p+; 10n0; 10e–) Figure 2.5a

  26. Chemically Reactive Elements • Outermost energy level not fully occupied by electrons • Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability

  27. (b) Chemically reactive elements Outermost energy level (valence shell) incomplete 4e 2e 1e Hydrogen (H) (1p+; 0n0; 1e–) Carbon (C) (6p+; 6n0; 6e–) 1e 6e 8e 2e 2e Oxygen (O) (8p+; 8n0; 8e–) Sodium (Na) (11p+; 12n0; 11e–) Figure 2.5b

  28. Types of Chemical Bonds • Ionic • Covalent • Hydrogen

  29. Ionic Bonds • Ions are formed by transfer of valence shell electrons between atoms • Anions (– charge) have gained one or more electrons • Cations (+ charge) have lost one or more electrons • Attraction of opposite charges results in an ionic bond

  30. + – Sodium atom (Na) (11p+; 12n0; 11e–) Chlorine atom (Cl) (17p+; 18n0; 17e–) Sodium ion (Na+) Chloride ion (Cl–) Sodium chloride (NaCl) (a) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron. (b) After electron transfer, the oppositely charged ions formed attract each other. Figure 2.6a-b

  31. Formation of an Ionic Bond • Ionic compounds form crystals instead of individual molecules • NaCl (sodium chloride)

  32. CI– Na+ (c) Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals. Figure 2.6c

  33. Covalent Bonds • Formed by sharing of two or more valence shell electrons • Allows each atom to fill its valence shell at least part of the time

  34. Reacting atoms Resulting molecules + or Structural formula shows single bonds. Molecule of methane gas (CH4) Hydrogen atoms Carbon atom (a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms. Figure 2.7a

  35. Reacting atoms Resulting molecules + or Structural formula shows double bond. Molecule of oxygen gas (O2) Oxygen atom Oxygen atom (b) Formation of a double covalent bond: Two oxygen atoms share two electron pairs. Figure 2.7b

  36. Reacting atoms Resulting molecules + or Structural formula shows triple bond. Molecule of nitrogen gas (N2) Nitrogen atom Nitrogen atom (c) Formation of a triple covalent bond: Two nitrogen atoms share three electron pairs. Figure 2.7c

  37. Covalent Bonds • Sharing of electrons may be equal or unequal • Equal sharing produces electrically balanced nonpolar molecules • CO2

  38. Figure 2.8a

  39. Covalent Bonds • Unequal sharing by atoms with different electron-attracting abilities produces polar molecules • H2O • Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen • Atoms with one or two valence shell electrons are electropositive, e.g., sodium

  40. Figure 2.8b

  41. Figure 2.9

  42. Hydrogen Bonds • Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule • Common between dipoles such as water • Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape PLAY Animation: Hydrogen Bonds

  43. + – Hydrogen bond (indicated by dotted line) + + – – – + + + – (a) The slightly positive ends (+) of the watermolecules become aligned with the slightlynegative ends (–) of other water molecules. Figure 2.10a

  44. (b) A water strider can walk on a pond because of the highsurface tension of water, a result of the combinedstrength of its hydrogen bonds. Figure 2.10b

  45. Chemical Reactions • Occur when chemical bonds are formed, rearranged, or broken • Represented as chemical equations • Chemical equations contain: • Molecular formula for each reactant and product • Relative amounts of reactants and products, which should balance

  46. Examples of Chemical Equations H + H  H2 (hydrogen gas) 4H + C  CH4 (methane) (reactants) (product)

  47. Patterns of Chemical Reactions • Synthesis (combination) reactions • Decomposition reactions • Exchange reactions

  48. Synthesis Reactions • A + B  AB • Always involve bond formation • Anabolic

  49. (a) Synthesis reactions Smaller particles are bonded together to form larger, more complex molecules. Example Amino acids are joined together to form a protein molecule. Amino acid molecules Protein molecule Figure 2.11a

  50. Decomposition Reactions • AB  A + B • Reverse synthesis reactions • Involve breaking of bonds • Catabolic

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