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Biology Biochemistry Unit Chapter 2 The Chemistry of Life PowerPoint Presentation
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Biology Biochemistry Unit Chapter 2 The Chemistry of Life

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Biology Biochemistry Unit Chapter 2 The Chemistry of Life

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Biology Biochemistry Unit Chapter 2 The Chemistry of Life

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  1. Biology Biochemistry Unit Chapter 2 The Chemistry of Life

  2. Topics • Atoms, Elements & Molecules • Chemical Bonds • Chemical Reactions • Mixtures & Solutions • Acids & Bases • Water • Cell Transport • Life Substances • Proteins • Carbohydrates • Nucleic acids • Lipids

  3. The Atom – makes up all matter • The smallest particle of an element • Made of 3 “subatomic particles” Nucleus is made of: positive PROTONS (p+) and neutral NEUTRONS (n0) Forming a cloud around the nucleus are negative electrons (e-) Figure 6.2

  4. Elements • Cannot be broken down into a simpler substance • 90 occur naturally; the rest are synthetic or radioactive • # of Protons gives characteristic nature – state, reactivity, etc • Many are needed by Living Orgs. for 8 characteristics….

  5. Periodic Table of Elements

  6. 96% Living things = CHNOWe’re more than CHNOPS Look at Table 6.1

  7. Elements: ~25 elements are essential to life What are trace elements? Found in the body in very small amounts, yet play vital roles!

  8. Elements All atoms are the same Represented by chemical symbol Compounds Combinations of different amounts of different atoms Represented by Chemical formulas Elements & Compounds

  9. Elements, Molecules & Compounds • ELEMENTS: • Symbols can be one capital letter or a capital letter with a lower case letter • For example: C = carbon; Ca = Calcium • MOLECULES • The combination of 2 or more elements H2O; O2 • COMPOUNDS: • A substance made of many different elements bonded together • Example: H2O = ? • Made of 2 elements – hydrogen and oxygen • NOTE: the subscripts tell # ofatoms of each element you have

  10. On your notes sheet Let’s try some element identification and atom counting…

  11. How many atoms?Try These… • USE APPENDIX D – The Periodic Table • KF • Potassium = __; Fluorine = __ • NaBr • Sodium = __; Bromine = __ • H3PO4 • ????? = __; ?????= __;?????= __ • CuSO4 • ?????? = __; ?????? = __; ?????? = __ • Pb(NO3)2 • ?????? = __; ?????? = __; ?????? = __

  12. How many atoms?Try These… • USE APPENDIX D – The Periodic Table • KF • Potassium = 1; Fluorine = 1 • NaBr • Sodium = 1; Bromine = 1 • H3PO4 • Hydrogen = 3; Phosphorus = 1; Oxygen = 4 • CuSO4 • Copper = 1; Sulfur =1; Oxygen = 4 • Pb(NO3)2 • Lead = 1; Nitrogen = 2; Oxygen = 6

  13. Parts of the atom:Protons, neutrons, electrons • The atomic number tells how many Protons • Neutral atoms have = number of P & E • Example: Carbon = element 6 • Has 6 protons and 6 electrons • What about n0? • n0 number = atomic mass (rounded) – # protons • So for carbon: 12 – 6 = 6 n0 Atomic # Atomic MASS

  14. On your notes sheet again…

  15. Parts of the atom:Try these… • NOTE: round the atomic mass • Calcium • Protons = __; Neutrons __; Electrons __ • Nickel • Protons = __; Neutrons __; Electrons __ • Gold • Protons = __; Neutrons __; Electrons __

  16. Parts of the atom:Try these…How many did you get correct? • NOTE: round the atomic mass • Calcium • Protons = 20; Neutrons 20; Electrons 20 • Nickel • Protons = 28; Neutrons 31; Electrons 28 • Gold • Protons = 79; Neutrons 118; Electrons 79

  17. Atoms that are different: ISOTOPES • Neutron (n0 ) # can vary • Atoms of the same element with different # neutrons = isotopes • Note: the p+ # NEVER CHANGES!!!

  18. Atoms that are different:Isotopes • Isotopes are represented by the number of neutrons • The number of p+ and e- stay the same • Carbon-12 = 6p+;6e-;6 n • Carbon-13 = 6 p+; 6e-;7n • Carbon-14 = ?????????????

  19. Ions – different # of electrons • Elements with more or less electrons than their atomic number designates • Br - = Bromine with one extra electron • Na + = Sodium with one less electron

  20. Study Guide • 1 - 3

  21. Electrons are constantly in motion around the nucleus • Electrons are attracted to the positive charges in the nucleus (protons) so they remain in orbit around the nucleus • Electrons move around in energy levels called electron clouds • Each can hold a certain number of electrons -first level can hold up to 2 e- -second level can hold up to 8 e- -third level can hold up to 8 e- - 4th level – 18e-

  22. The Atom:Electron clouds & “dot” diagrams

  23. Look at the periodic table to clear up any confusion… 2 8 8 18

  24. The Atom:Electron clouds • EXAMPLE: Carbon has 6 electrons • They are arranged as: • 2 in the first e- cloud • 4 in the second e- cloud

  25. THE BOTTOM LINE • For dot diagram purposes use the rule: • 2 • 8 • 8 • 18 • Let’s try some examples… What element is this?

  26. The Atom:Try these… • Which element do these atomic structures represent?

  27. Why is this important? • Shows the VALENCE e-’s (outer) • Determines an elements bonding potential • its REACTIVITY! • Example: • How many valence e-’s does oxygen have? • How many does it “want” to be stable?

  28. Chemical bonds:Forming a compound • Bond – to join together atoms using electron energy/force • Atoms are most stable when their outer e- cloud is full • Elements can share, steal or lose electrons to fill their outer (valence) e- cloud • This causes stability • Types of Bonds: Covalent, Polar Covalent, Ionic, Hydrogen & Vander Waals Forces

  29. 1. Covalent bonds EXAMPLE: Oxygen alone only has 6 electrons in its outer E level If it shares 2 more it would have 8 in its outer E level and be stable ”CO” = share Figure 6.6

  30. Covalent Bonds: Dogs of equal strength • Bone = Electron • Covalent bonds - two or more dogs with equal attraction to bones • Dogs (atoms) are identical, then the dogs share the pairs of available bones evenly • Since one dog does not have more of the bone than the other dog, the charge is evenly distributed among both dogs • The molecule is not "polar" meaning one side does not have more charge than the other.

  31. 2. Ionic bonds • When atoms give or take electrons; not sharing • The atoms with extra (or less) e- are now not neutral = ION = bond forms Figure 6.7

  32. Ionic Bonds: One big greedy thief dog! • One big greedy dog steeling other dog's bone • Bone represents the e- that is up for grabs • then when the big dog gains an e- he becomes negatively charged • the little dog who lost the e- becomes positively charged • The two ions (that's where the name ionic comes from) are attracted strongly to each other as a result of the opposite charges.

  33. Bond Strengths • How much energy is stored bond/How much energy it takes to break the bond 1. Triple Covalent 2. Double Covalent 3. Covalent Single 4. Ionic 5. Hydrogen*

  34. Think about it! • Carbon (bonded to H) can also bond to Nitrogen • How many electrons are shared between C& N? • What would this bond look like? • This type of bond is found in proteins such as muscle. Why?

  35. Chemical Reactions • Bonds of atoms are broken and re-formed into new substances • Reactions are written as “chemical equations” which show reactants and products • All reactions in an organism = Metabolism What does balanced mean?

  36. Single, Double and Triple Covalent Bonds • 1 electron shared = single bond • Ex: C - H • 2 electrons shared = double bond • Ex: O= O

  37. Chemical Reactions • EXAMPLE 6CO2 + 6H2O  C6H12O6 + 6O2 This is the chemical equation for the reaction that occurs…. Notice the COEFFICENTS WHAT IS THEIR PURPOSE?

  38. Kinesthetic Analogy • H + Cl HCl Hydrochloric Acid - Ionic • H + H  H2 Gas - Covalent • Na + Cl  NaCl Sodium Chloride - Ionic

  39. All the “biochemistry” in one of your cells every second…

  40. Study Guide • 3 & 5

  41. Properties of Water We’ll do water now

  42. WaterHydrophilic vs. Hydrophobic • Hydrophilic – substances have affinity water • Hydrophobic – substances that repel water; form clusters called hydrophobic interactions. • Ex: Oil and H2O • Like dissolves like

  43. Water is Polar • Many covalent bonds (like in water) don’t share the e- equally • One side of the molecule is slightly (–) and the other is slightly (+) POLAR MOLECULE (Polarity) Oxygen “hogs” the e- making it more -

  44. Hydrogen bondsThe polarity of water helps it to bond to itself…

  45. Properties of Water • Universal Solvent • Water resists temperature changes • Water expands when it freezes – less dense • Cohesion E. Surface Tension • Adhesion G. Exists in 3 states – solid, liquid, gas (vapor) H. Evaporative coolant I. Capillarity

  46. Since it has a charge (polar) & can form H-bonds, water molecules can surround OTHER molecules • What is this called?

  47. A. Water is the UNIVERSAL SOLVENT Here we see:A crystal of table salt dissolving in water • DISSOLVING!

  48. Water Properties A. Powerful, UNIVERSAL SOLVENT – due to polarity; dissolves many solutes • Polar charges attract/cause ion formation of solutes  dissolving (breaking apart) solutes

  49. B. Water resists temperature changes • Helps cells maintain homeostasis • Cells don’t change in temperature rapidly