Chapter 2 - Molecules of Life (Biochemistry)! Periodic Table of Elements " " Figure 2.2!

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1 Chapter 2 - Chemistry Chapter 2 - Molecules of Life (Biochemistry) Element = a quantity of matter composed of atoms of the same type Atoms: net charge = zero Protons (+, mass 1) Neutrons (0, mass 1) Electrons (-, mass negligible) Atom has same # of protons and electrons so charge = 0" Atomic number = # protons Mass number = "# protons + # neutrons 1 Periodic Table of Elements " " Figure Na 23 Element 8 O 16 Atomic number Mass number CHONPS, Na, K, Ca 2 1

2 Chapter 2 - Chemistry Electron Shells" " " " Figure 2.8 What is the charge on an atom of oxygen? Draw a diagram like the ones above for Na 11 and Cl Ions - The Octet Rule (Rule of Eight) Atoms can gain or lose electrons Except for the first electron shell, the outermost (valence) shell can hold 8 electrons (This applies to all atoms that you need to know about.) E.g. Sodium atom (Na ) loses one electron Na+ Giving something away is a positive thing to do Positively charge ion = cation (ca+ion) E.g. Chlorine (Cl ) atom gains one electron Cl- Negatively charged ion = anion 4 2

3 Chapter 2 - Chemistry Chemical Bonds Contain energy that is released when the bond is broken glucose (a sugar) CO 2 + H 2 O+ energy Energy is required to form chemical bonds CO 2 + H 2 O + energy glucose Catabolic? Anabolic? 5 Ionic Bonds " " " " " Figure 2.10 Anions (-) and cations (+) attract each other" 6 3

4 Chapter 2 - Chemistry Covalent Bonds Do not involve ions - Atoms share electrons Strong - hard to break (requires enzymes) Most of the energy we get from food involves breaking covalent bonds A B A + B + Energy glucose + oxygen carbon dioxide + water+ Energy C 6 H 12 O O 2 6 CO H 2 O + Energy 7 Single Covalent Bonds " " Figure 2.11 How many electrons are in each atom s outermost shell? (Charges on H 2 O are misleading. i.e., WRONG. See slide 11, 12.) Shorthand notation: a single line connecting two atoms indicates a pair of covalently shared electrons. O H H 8 4

5 Chapter 2 - Chemistry Double Covalent Bonds " " Figure 2.9b How many electrons are in each atom s outermost shell? Carbon dioxide: CO 2 Chemical notation: O=C=O, the double lines (=) indicate two pairs of shared electrons in CO 2 9 Covalent Bonds - 2 Nonpolar covalent bonds Electrons shared equally between atoms Previous slide for CO 2 showed nonpolar covalent bonds Polar covalent bonds Electrons not shared equally One atom hogs the electrons This leads to the formation of hydrogen bonds. 10 5

6 Chapter 2 - Chemistry Hydrogen Bonds (H-bonds) " Figure 2.11a Relatively weak bonds, but lots of them together can be strong. Result from unequal sharing of electrons in polar covalent molecules. Partial positive and negative charges on different molecules attract each other. Water is a polar covalent molecule. Electrons are shared unequally between atoms. Figure 2.10a is misleading. δ + 2δ - δ + 11 Correction to Figure 2.11a 2δ - δ + δ + The charges are not full charges like those on a sodium or chloride ion. They are partial charges, represented by a δ (delta). 12 6

7 Chapter 2 - Chemistry Hydrogen Bonding " " " (Martini 2006) 13 Properties of Water A polar molecule - will form hydrogen bonds with itself 1. Water has a high heat capacity - hydrogen bonds again. Absorbs heat Stabilizes temperature Evaporation removes heat 2. Water is a good lubricant - little friction between individual water molecules. 14 7

8 Chapter 2 - Chemistry Properties of Water - 2 " " (Martini 2006) 3. Water is a good solvent. Solute Solvent Forms spheres of hydration 15 Properties of Water Ice floats (Learn something new every day.) Freezing stabilizes H-bonds - makes ice less dense than water. 5. Water is cohesive. Water molecules stick together. H-bonds again. 6. Water has surface tension. H-bonds again. 16 8

9 Chapter 2 - Chemistry Acids and Bases - ph (not Ph, ph, or PH) Acids give up H + (hydrogen ions) "Strong acids completely dissociate (fall apart): E.g. HCl H + + Cl - "Weak acids do not completely dissociate: E.g. H 2 CO 3 H + + HCO 3- Bases take up H + or lose OH - E.g. NaOH Na + + OH - 17 ph is a Measure of Hydrogen Ion Concentration ph = -log[h + ] "Examples of log: log(100) = log(10 2 ) = 2 log(0.01) = log(10-2 ) =- 2 Pure water at 25 C has [H + ] of 10-7 moles/liter " " "ph = -log[10-7 ] " " " = -(-7) " " " = 7 ph of 7 is said to be neutral 18 9

10 Chapter 2 - Chemistry ph Scale" " " " Blood: Gastric juice: Vaginal fluids: Semen: Urine: Why does it make sense that all ph values listed above aren t the same as that of blood? 19 Buffers Buffers help maintain ph within desirable limits." Composed of a weak acid and a weak base "Carbonic acid <----> Hydrogen ion + bicarbonate ion " (a weak acid) " " " (a weak base) Chemically: " "H 2 CO 3 < > H + + HCO - 3 Add acid (HCl), reaction moves to left (Why?) Add base (NaOH), reaction moves to right (Why?) Both prevent large changes in [H + ] 20 10

11 Chapter 2 - Chemistry Types of Chemical Reactions 1. Dehydration synthesis Make more complex molecules from simpler ones Require input of energy A.K.A. endergonic or anabolic reactions Water is formed (dehydration occurs) A-OH + B-H + energy > A-B + H 2 O e.g. synthesis of glycogen from glucose Dehydration Synthesis " Figure 2.15a Water leaves: dehydration 22 11

12 Chapter 2 - Chemistry 2. Hydrolysis Reaction A.K.A. decomposition reaction Break larger molecule into smaller parts Water is added (hydrolysis means to break with water ) Energy is released (exergonic or catabolic) A-B + H 2 O > A-H + B-OH + energy e.g. hydrolysis of glycogen to glucose Hydrolysis Reaction " " " Figure 2.15b Water is used to break this bond: hydrolysis 24 12

13 Chapter 2 - Chemistry Many (most) Biological Reactions are Reversible 3. Reversible reactions Substrates and products often exist in equilibrium Generically: AB A + B If the concentration of B goes up, which way does the reaction go? e.g. H 2 CO 3 H + + HCO 3-25 Organic Molecules - Carbohydrates General formula = (CH 2 O) n I.e. carbon and water (carbo- + hydro-) General functions: Energy source and energy reserve " " "Glucose and glycogen Structural molecules " " "Deoxyribose in DNA backbone e.g. "Monosaccharides (e.g. glucose, fructose) " "Disaccharides (e.g. sucrose, maltose) " "Polysaccharides (e.g. glycogen) 26 13

14 Chapter 2 - Chemistry Carbohydrates - 2 " " " " Figure 2.16 For glucose n = 6, so (CH 2 O) n = C 6 H 12 O 6 Different ways to represent glucose: 27 Carbohydrates - Polysaccharides Glycogen A string of glucose units Stored for energy in animal cells Starch A string of glucose units Stored for energy in plant cells Cellulose A string of glucose units A structural molecule in plants Why can t we digest cellulose in trees? 28 14

15 Chapter 2 - Chemistry Lipids (Fats) Not water soluble" Some types: Triglycerides (saturated and unsaturated fats) Phospholipids Steroids 29 Lipids - Triglycerides General formula: Glycerol + 3 fatty acids Functions: Energy storage (twice as much energy per gram as carbohydrates) Energy source (fatty acids) Types: Saturated fats (solids at room temp.) Unsaturated fats (liquids at room temp.) 30 15

16 Chapter 2 - Chemistry Lipids - Triglycerides " " " Figure 2.19 Glycerol Fatty acid Unsaturated Saturated 31 Lipids - Phospholipids General formula: Two fatty acid tails + phosphatecontaining head General functions: Important component of cell membranes Are polar molecules Hydrophilic head Hydrophobic tail 32 16

17 Chapter 2 - Chemistry Lipids - Phospholipids " " " Figure _ Note electrical charges More about this in Ch Lipids Steroids " " " " Figure 2.21 General structures: General functions: Hormones A component of cell membranes Cholesterol Estrogen Testosterone 34 17

18 Chapter 2 - Chemistry Proteins General formula: Chain of amino acids joined by peptide (covalent) bonds" General functions: Structural function - Muscle, bone, hair Enzymes - Chemical catalysts Oxygen carriers (e.g. hemoglobin, myoglobin) Hormones (e.g. insulin, growth hormone) 35 Amino Acids 1 " " " " Figure 2.22 R group or side chain determines the amino acid s properties. Hydrophilic? Hydrophobic? What do you think R stands for? (Acid) 36 18

19 Chapter 2 - Chemistry Protein Structure Summary " " Figure 2.24 Primary: Linear sequence of amino acids Tertiary: Overall 3-D shape of a polypeptide Secondary: Local folding into helices, sheets Quaternary: Interactions between two or more polypeptides 37 Nucleic Acids General formula: Composed of nucleotides" Sugar Nitrogen-containing base Phosphate DNA - double stranded RNA - single stranded ATP 38 19

20 Chapter 2 - Chemistry Nucleotide Structure " " " Figure 2.26 Pentose sugar: "Ribose in RNA and ATP " " "Deoxyribose in DNA 39 Nitrogenous Bases A = Adenine T = Thymine C = Cytosine G = Guanine DNA base pairing rules: A is complementary to T G is complementary to C RNA base pairing rules: A is complementary to U G is complementary to C Why are DNA and RNA different? Because nature does it this way. Pretty simple, eh? 40 20

21 Chapter 2 - Chemistry DNA Structure" " " " " Figure To Be Covered Later More about Nucleic acids will be covered in Chapter 3 and 21. We will cover Enzyme Function (pp ) with Chapter 3 material

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