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Chemistry Of Life

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  1. Chemistry Of Life

  2. Elemental Life

  3. Primary Elemental Life - SPONCH

  4. Primary Elemental Life - SPONCH

  5. Atom Basics • Relevant subatomic particles include • Neutrons, which have no electrical charge • Protons, which are positively charged • Electrons, which are negatively charged

  6. Bonding – its all about making atoms happy Ionic – give and take Covalent – Share e- Based on Electronegativity

  7. Ionic Bonds • In some cases, atoms strip electrons away from their bonding partners

  8. Electron transfer between two atoms creates ions • Ions • Are atoms with more or fewer electrons than usual • Are charged atoms

  9. The lone valence electron of a sodium atom is transferred to join the 7 valence electrons of a chlorine atom. Each resulting ion has a completed valence shell. An ionic bond can form between the oppositely charged ions. – + 1 2 Cl Na Na Cl Cl– Chloride ion (an anion) Na+ Sodium on (a cation) Na Sodium atom (an uncharged atom) Cl Chlorine atom (an uncharged atom) Sodium chloride (NaCl) Ionic Bond – Anions and Cations Figure 2.13

  10. Hydrogen atoms (2 H) In each hydrogen atom, the single electron is held in its orbital by its attraction to the proton in the nucleus. + + 1 2 3 When two hydrogen atoms approach each other, the electron of each atom is also attracted to the proton in the other nucleus. + + The two electrons become shared in a covalent bond, forming an H2 molecule. + + Hydrogen molecule (H2) Formation of Covalent bond Figure 2.10

  11. In a nonpolar covalent bond • The atoms have similar electronegativities • Share the electron equally

  12. Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen. d– This results in a partial negative charge on the oxygen and a partial positive charge on the hydrogens. O H H d+ d+ H2O Polar covalent bond • atoms have differing electronegativitiesShare the electrons unequally Figure 2.12

  13. H Water (H2O) O A hydrogen bond results from the attraction between the partial positive charge on the hydrogen atom of water and the partial negative charge on the nitrogen atom of ammonia. H  +  – Ammonia (NH3) N H H d+ + H Figure 2.15 Hydrogen Bonds • Forms when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom  –  +  +

  14. Van der Waals Interactions Occur when transiently positive and negative regions of molecules attract each other

  15. Weak chemical bonds • Reinforce the shapes of large molecules • Help molecules adhere to each other

  16. – Hydrogenbonds + H – + H + –  – + Figure 3.2 • The polarity of water molecules • Allows them to form hydrogen bonds with each other • Contributes to the various properties water exhibits

  17. 5 Properties of Water

  18. 5 Properties of Water 1. Cohesion 2. Adhesion 3. Less Dense as solid 4. Solvent 5. High Spec. Heat

  19. 5 Properties of Water 1. Cohesion A. Evap Cooling 2. Adhesion B. Sticks to self 3. Less Dense C. Ice floats as solid D. Sticks to other 4. Solvent E. Metabolism 5. High Spec. Heat F. Climate Control G. Transpiration H. Health of Climate Varying Aquatic Biomes

  20. 5 Properties of Water 1. Cohesion A. Evap Cooling 2. AdhesionB. Sticks to self 3. Less DenseC. Ice floats as solid D. Sticks to other 4. Solvent E. Metabolism 5. High Spec. Heat F. Climate Control G. Transpiration H. Health of Climate Varying Aquatic Biomes

  21. Cohesion • Water molecules exhibit cohesion • Cohesion • Is the bonding of a high percentage of the molecules to neighboring molecules • Is due to hydrogen bonding

  22. Water conducting cells 100 µm Figure 3.3 Cohesion • Helps pull water up through the microscopic vessels of plants

  23. Figure 3.4 Surface tension • Is a measure of how hard it is to break the surface of a liquid

  24. Water’s High Specific Heat • The specific heat of a substance • Is the amount of heat that must be absorbed or lost for 1 gram of that substance to change its temperature by 1ºC

  25. Moderation of Temperature • Water moderates air temperature • By absorbing heat from air that is warmer and releasing the stored heat to air that is cooler

  26. Evaporative Cooling • Evaporation • Is the transformation of a substance from a liquid to a gas

  27. Evaporative cooling • Is due to water’s high heat of vaporization • Allows water to cool a surface

  28. Insulation of Bodies of Water by Floating Ice • Solid water, or ice • Is less dense than liquid water • Floats in liquid water

  29. Hydrogen bond Liquid water Hydrogen bonds constantly break and re-form Ice Hydrogen bonds are stable Figure 3.5 • The hydrogen bonds in ice • Are more “ordered” than in liquid water, making ice less dense

  30. The Solvent of Life • Water is a versatile solvent due to its polarity • It can form aqueous solutions

  31. Negative oxygen regions of polar water molecules are attracted to sodium cations (Na+). – Na+ + + – + – – Positive hydrogen regions of water molecules cling to chloride anions (Cl–). Na+ – + + Cl – Cl– + – – + – + – – Figure 3.6 The different regions of the polar water molecule can interact with ionic compounds called solutes and dissolve them

  32. This oxygen is attracted to a slight positive charge on the lysozyme molecule. – + This oxygen is attracted to a slight negative charge on the lysozyme molecule. (b) Lysozyme molecule (purple) in an aqueous environment such as tears or saliva (a) Lysozyme molecule in a nonaqueous environment (c) Ionic and polar regions on the protein’s Surface attract water molecules. Figure 3.7 Water can also interact with polar molecules such as proteins

  33. Solute Concentration in Aqueous Solutions • Since most biochemical reactions occur in water • It is important to learn to calculate the concentration of solutes in an aqueous solution

  34. Acids and Bases • An acid • Is any substance that increases the hydrogen ion concentration of a solution • A base • Is any substance that reduces the hydrogen ion concentration of a solution

  35. pH Scale 0 1 Battery acid 2 Digestive (stomach) juice, lemon juice Vinegar, beer, wine, cola 3 Increasingly Acidic [H+] > [OH–] 4 Tomato juice 5 Black coffee Rainwater 6 Urine Neutral [H+] = [OH–] 7 Pure water Human blood 8 Seawater 9 10 Increasingly Basic [H+] < [OH–] Milk of magnesia 11 Household ammonia 12 Household bleach 13 Oven cleaner 14 Figure 3.8 pH Scale

  36. Buffers • The internal pH of most living cells • Must remain close to pH 7 • Are substances that minimize changes in the concentrations of hydrogen and hydroxide ions in a solution • Consist of an acid-base pair that reversibly combines with hydrogen ions

  37. The Threat of Acid Precipitation • Acid precipitation • Refers to rain, snow, or fog with a pH lower than pH 5.6 • Is caused primarily by the mixing of different pollutants with water in the air

  38. 0 1 Moreacidic 2 3 Acidrain 4 5 Normalrain 6 7 8 9 10 11 12 13 Morebasic 14 Figure 3.9 Acid precipitation damage

  39. Biological Molecules • Based on Carbon • Four possible bonds • Many isomers

  40. H H H C H H C H H H H H H H (a) Structural isomers H C C C C C H H C H C C H H H H H H H H H X X X C C C C (b) Geometric isomers X H H H CO2H CO2H C C (c) Enantiomers H H NH2 NH2 CH3 CH3 Figure 4.7 A-C Three types of isomers Structural Geometric Enantiomers

  41. L-Dopa (effective against Parkinson’s disease) D-Dopa (biologically inactive) Figure 4.8 Enantiomers

  42. The Functional Groups Most Important in the Chemistry of Life • Functional groups • Are the chemically reactive groups of atoms within an organic molecule

  43. OH CH3 Estradiol HO Female lion OH CH3 CH3 O Testosterone Male lion Figure 4.9

  44. Functional Groups

  45. Functional Groups

  46. 1 HO H 3 2 H HO Unlinked monomer Short polymer Dehydration removes a watermolecule, forming a new bond H2O 1 2 3 4 HO H Longer polymer (a) Dehydration reaction in the synthesis of a polymer Figure 5.2A The Synthesis and Breakdown of Polymers • Dehydration Synthesis

  47. 1 3 HO 4 2 H Hydrolysis adds a watermolecule, breaking a bond H2O 1 2 H HO 3 H HO (b) Hydrolysis of a polymer Figure 5.2B Hydrolysis

  48. Sugars • Monosaccharides • Are the simplest sugars • Can be used for fuel • Can be converted into other organic molecules • Can be combined into polymers

  49. Triose sugars(C3H6O3) Pentose sugars(C5H10O5) Hexose sugars(C6H12O6) H H H H O O O O C C C C H C OH H C OH H C OH H C OH H C OH H C OH HO C H HO C H Aldoses H H C OH H C OH HO C H H C OH H C OH H C OH Glyceraldehyde H C OH H C OH H Ribose H H Glucose Galactose H H H H C OH H C OH H C OH C O C O C O HO C H H C OH H C OH Ketoses H C OH H C OH H Dihydroxyacetone H C OH H C OH H C OH H Ribulose H Figure 5.3 Fructose Examples of monosaccharides

  50. Disaccharides • Consist of two monosaccharides • Are joined by a glycosidic linkage