1. Copyright © 2010 Pearson Education, Inc. Biochemistry Part A Biochemistry- the chemistry of living things

2. Copyright © 2010 Pearson Education, Inc. Matter Anything that has mass and occupies space States of matter: 1.Solid—definite shape and volume 2.Liquid—definite volume, changeable shape 3.Gas—changeable shape and volume

3. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Composition of Matter Atoms Unique building blocks for each element Atomic symbol: one- or two-letter chemical shorthand for each element

6. Copyright © 2010 Pearson Education, Inc. Major Elements of the Human Body Oxygen (O) Carbon (C) Hydrogen (H) Nitrogen (N) About 96% of body mass

7. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Trace Elements of the Human Body < 0.01% of body mass: Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)

9. Copyright © 2010 Pearson Education, Inc. Atomic Structure Determined by numbers of subatomic particles Nucleus consists of neutrons and protons

10. Copyright © 2010 Pearson Education, Inc. Atomic Structure Neutrons No charge Mass = 1 atomic mass unit (amu) Protons Positive charge Mass = 1 amu

11. Copyright © 2010 Pearson Education, Inc. Atomic Structure Electrons Orbit nucleus Equal in number to protons in atom Negative charge 1/2000 the mass of a proton (0 amu)

12. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Models of the Atom Planetary model—oversimplified, outdated model Incorrectly depicts fixed circular electron paths Useful for illustrations (as in the text)

14. Copyright © 2010 Pearson Education, Inc. Figure 2.1 (a) Planetary model(b) Orbital model Helium atom 2 protons (p + ) 2 neutrons (n 0 ) 2 electrons (e – ) Helium atom 2 protons (p + ) 2 neutrons (n 0 ) 2 electrons (e – ) Nucleus ProtonNeutronElectron cloud Electron

15. Copyright © 2010 Pearson Education, Inc. Identifying Elements Atoms of different elements contain different numbers of subatomic particles Compare hydrogen, helium and lithium (next slide)

16. Copyright © 2010 Pearson Education, Inc. Figure 2.2 Proton Neutron Electron Helium (He) (2p + ; 2n 0 ; 2e – ) Lithium (Li) (3p + ; 4n 0 ; 3e – ) Hydrogen (H) (1p + ; 0n 0 ; 1e – )

17. Copyright © 2010 Pearson Education, Inc. Identifying Elements Atomic number = number of protons in nucleus

18. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Identifying Elements Atomic weight = average of mass numbers of all isotopes

20. Copyright © 2010 Pearson Education, Inc. Figure 2.3 Proton Neutron Electron Deuterium ( 2 H) (1p + ; 1n 0 ; 1e – ) Tritium ( 3 H) (1p + ; 2n 0 ; 1e – ) Hydrogen ( 1 H) (1p + ; 0n 0 ; 1e – )

21. Copyright © 2010 Pearson Education, Inc. Radioisotopes Spontaneous decay (radioactivity) Similar chemistry to stable isotopes Can be detected with scanners

22. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Molecules and Compounds Most atoms combine chemically with other atoms to form molecules and compounds Molecule—two or more atoms bonded together (e.g., H 2 or C 6 H 12 O 6 ) Compound—two or more different kinds of atoms bonded together (e.g., C 6 H 12 O 6 )

24. Copyright © 2010 Pearson Education, Inc. Chemically Inert Elements Stable and unreactive Outermost energy level fully occupied or contains eight electrons

25. Copyright © 2010 Pearson Education, Inc. Figure 2.5a Helium (He) (2p + ; 2n 0 ; 2e – ) Neon (Ne) (10p + ; 10n 0 ; 10e – ) 2e 8e (a) Chemically inert elements Outermost energy level (valence shell) complete

26. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Figure 2.5b 2e 4e 2e 8e 1e (b) Chemically reactive elements Outermost energy level (valence shell) incomplete Hydrogen (H) (1p + ; 0n 0 ; 1e – ) Carbon (C) (6p + ; 6n 0 ; 6e – ) 1e Oxygen (O) (8p + ; 8n 0 ; 8e – ) Sodium (Na) (11p + ; 12n 0 ; 11e – ) 2e 6e

28. Copyright © 2010 Pearson Education, Inc. Types of Chemical Bonds Ionic Covalent Hydrogen

29. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Figure 2.6a-b Sodium atom (Na) (11p + ; 12n 0 ; 11e – ) Chlorine atom (Cl) (17p + ; 18n 0 ; 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.

31. Copyright © 2010 Pearson Education, Inc. Formation of an Ionic Bond Ionic compounds form crystals instead of individual molecules NaCl (sodium chloride)

32. Copyright © 2010 Pearson Education, Inc. Figure 2.6c CI – Na + (c) Large numbers of Na + and Cl – ions associate to form salt (NaCl) crystals.

33. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Figure 2.7a + Hydrogen atoms Carbon atom Molecule of methane gas (CH 4 ) Structural formula shows single bonds. (a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms. or Resulting moleculesReacting atoms

35. Copyright © 2010 Pearson Education, Inc. Figure 2.7b or Oxygen atom Oxygen atom Molecule of oxygen gas (O 2 ) Structural formula shows double bond. (b) Formation of a double covalent bond: Two oxygen atoms share two electron pairs. Resulting moleculesReacting atoms +

36. Copyright © 2010 Pearson Education, Inc. Figure 2.7c + or Nitrogen atom Nitrogen atom Molecule of nitrogen gas (N 2 ) Structural formula shows triple bond. (c) Formation of a triple covalent bond: Two nitrogen atoms share three electron pairs. Resulting moleculesReacting atoms

37. Copyright © 2010 Pearson Education, Inc. Covalent Bonds Sharing of electrons may be equal or unequal Equal sharing produces electrically balanced nonpolar molecules CO 2

38. Copyright © 2010 Pearson Education, Inc. Figure 2.8a

39. Copyright © 2010 Pearson Education, Inc. Covalent Bonds Unequal sharing by atoms with different electron-attracting abilities produces polar molecules H 2 O 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. Copyright © 2010 Pearson Education, Inc. Figure 2.8b

41. Copyright © 2010 Pearson Education, Inc. Figure 2.9

42. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. (a) The slightly positive ends (  + ) of the water molecules become aligned with the slightly negative ends (  – ) of other water molecules. ++ –– –– –– –– –– ++ ++ ++ ++ ++ Hydrogen bond (indicated by dotted line) Figure 2.10a

44. Copyright © 2010 Pearson Education, Inc. Figure 2.10b (b) A water strider can walk on a pond because of the high surface tension of water, a result of the combined strength of its hydrogen bonds.

45. Copyright © 2010 Pearson Education, Inc. 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. Copyright © 2010 Pearson Education, Inc. Examples of Chemical Equations H + H  H 2 (hydrogen gas) 4H + C  CH 4 (methane) (reactants)(product)

47. Copyright © 2010 Pearson Education, Inc. Patterns of Chemical Reactions Synthesis (combination) reactions Decomposition reactions Exchange reactions

48. Copyright © 2010 Pearson Education, Inc. Synthesis Reactions A + B  AB Always involve bond formation Anabolic

49. Copyright © 2010 Pearson Education, Inc. Figure 2.11a Example Amino acids are joined together to form a protein molecule. (a) Synthesis reactions Smaller particles are bonded together to form larger, more complex molecules. Amino acid molecules Protein molecule

50. Copyright © 2010 Pearson Education, Inc. Decomposition Reactions AB  A + B Reverse synthesis reactions Involve breaking of bonds Catabolic

51. Copyright © 2010 Pearson Education, Inc. Figure 2.11b Example Glycogen is broken down to release glucose units. Bonds are broken in larger molecules, resulting in smaller, less complex molecules. (b) Decomposition reactions Glucose molecules Glycogen

52. Copyright © 2010 Pearson Education, Inc. Chemical Reactions All chemical reactions are either exergonic or endergonic Exergonic reactions—release energy Catabolic reactions Endergonic reactions—products contain more potential energy than did reactants Anabolic reactions

53. Copyright © 2010 Pearson Education, Inc. Chemical Reactions All chemical reactions are theoretically reversible A + B  AB AB  A + B Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant Many biological reactions are essentially irreversible due to Energy requirements Removal of products

54. Copyright © 2010 Pearson Education, Inc. Rate of Chemical Reactions Rate of reaction is influenced by:  temperature   rate  particle size   rate  concentration of reactant   rate Catalysts:  rate without being chemically changed Enzymes are biological catalysts