CHAPTER 02 CHEMICAL BASIS OF LIFE

BIO 211: ANATOMY & PHYSIOLOGY I 1 CHAPTER 02 CHEMICAL BASIS OF LIFE Dr. Lawrence G. G. Altman www.lawrencegaltman.com Some illustrations are courtesy of McGraw-Hill. Some illustrations are courtesy of McGraw-Hill. I. Matter consists of chemical elements in pure form and in combinations called compounds Chemistry is fundamental to an understanding of life, because living organisms are made of matter. Matter: Anything that takes up space and has mass. Mass: A measure of the amount of matter an object contains. There is a slight difference between mass and weight. Mass: is the measure of the amount of matter an object contains; it stays the same regardless of changes in the object's position. Weight: is the measure of how strongly an object is pulled by earth's gravity, and it varies with distance from the earth's center. 2 The important point is that the mass of a body does not vary with its position, whereas weight does. So, for all practical purposes - as long as we are earthbound - weight can be used as a measure of mass.

3 II. Life requires about 25 chemical elements Element: A substance that cannot be broken down into other substances by chemical reactions. * All matter made of elements. * 92 naturally occurring. * Designated by a symbol of one/two letters. About 25 of the 92 naturally - occurring elements are essential to life. Biologically important elements include those shown on the right. 4 Trace element: Element required by an organism in extremely minute quantities. Though in small quantity, are indispensable for life. For example: B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V and Zn. Ultratrace element: Element required by an organism in extremely minute quantities but TOXIC at high levels. For example: Arsenic (Ar)

5 6 II. Life requires about 25 chemical elements (cont.) Elements can exist in combinations called compounds. Compound: A pure substance composed of two or more DIFFERENT elements combined in a fixed ratio. For example: NaCl (sodium chloride).

7 III. Atomic structure determines the behavior of an element Atom: Smallest possible unit of matter that retains (keeps) the physical and chemical properties of its element. Atoms of the same element share similar chemical properties. Made up of subatomic particles. A. Subatomic Particles The three most stable subatomic particles are: 1. Neutrons no charge (neutral) 2. Protons +1 electrostatic charge 3. Electrons - 1 electrostatic charge NEUTRON PROTON n 0 H + no charge +1 charge Found together in a dense core called the atomic nucleus. Nucleus is + (due to + protons) 1.009 dalton 1.007 dalton Masses of both are about the same (approx. 1 dalton) ELECTRON e - e - -1 charge Orbit around the nucleus; held together by electrostatic attraction to + charged nucleus 1/2000 dalton Mass is so small, usually not used to calculate atomic mass 8 NOTE: The Dalton is used to express mass at the atomic level. If an atom is electrically neutral, the number of protons EQUALS the number of electrons, which yields an electrostatically balanced charge.

9 III. Atomic structure determines the behavior of an element (cont.) B. Atomic Number and Atomic Mass Atomic number: Number of PROTONS in an atom of a particular element. All atoms of an element have the same atomic number. Written as a subscript to the left of the element's symbol. Example: 11 Na In a neutral atom, # protons = # electrons. 10 III. Atomic structure determines the behavior of an element (cont.) B. Atomic Number and Atomic Mass Mass number: Number of PROTONS + NEUTRONS in an atom. Each has a mass of approx 1 dalton. Written as a superscript to the left of an element s symbol. Example: 23 Na Is the approximate mass of the whole atom, since the mass of a proton and the mass of a neutron are each equal to 1 dalton. dalton: a measurement of mass at the atomic level.

11 III. Atomic structure determines the behavior of an element (cont.) B. Atomic Number and Atomic Mass Mass number: Can deduce the number of NEUTRONS by subtracting atomic number from atomic mass. # Neutrons = Mass number - Atomic number N = (protons + N) - (protons) The number of Protons (the Atomic number??) in an element is ALWAYS constant (the same).. The number of Neutrons in an element CAN vary. Now, try a few problems. 12 Protons + Neutrons MASS NUMBER ATOMIC NUMBER Protons only 23 11 Na 12 6 C # of electrons # of protons # of neutrons # of electrons # of protons # of neutrons

13 14 III. Atomic structure determines the behavior of an element (cont.) C. Isotopes Isotopes: Atoms of an element that have the same atomic number but different mass number. Have the same number of protons, but different number of neutrons. Atomic WEIGHT: Weighted average of the mixture Under natural conditions, elements occur as mixtures of isotopes. Different isotopes of the same element react chemically in the same way. Some isotopes are radioactive.

15 III. Atomic structure determines the behavior of an element (cont.) C. Isotopes Radioactive isotope: Unstable isotope in which the nucleus spontaneously decays emitting subatomic particles and/or energy as radioactivity. Loss of nuclear particles may transform (change) one element to another!! Has a fixed half - life Half life: Time for 50% of radioactive atoms in a sample to decay. 16

17 III. Atomic structure determines the behavior of an element (cont.) C. Isotopes Biological applications of radioactive isotopes include: 1. Dating geological strata (layers) and fossils. 2. Radioactive tracers Chemicals labeled with radioactive isotopes are used to trace the steps of a biochemical reaction or to determine the location of a particular substance within an organism. Radioactive isotopes are useful as biochemical tracers because they chemically react like the stable isotopes and are easily detected at low concentrations. Isotopes of P, N and H were used to determine DNA structure. Why? 18 III. Atomic structure determines the behavior of an element (cont.) C. Isotopes Biological applications of radioactive isotopes include: 2. Radioactive tracers (cont.) D. Energy Levels Used to diagnose disease (e.g. PET scanner). Because radioactivity can damage cell molecules, radioactive isotopes can also be hazardous. 3. Treatment of Cancer Example: radioactive cobalt. Electrons: Light negatively - charged particles that orbit around a nucleus. Are the only subatomic particles which are directly involved in chemical reactions. Cont. >>>

III. Atomic structure determines the behavior of an element (cont.) D. Energy Levels (cont.) Energy: Ability to do work. Potential Energy: Energy that matter stores because of its position or location. E. Electron Configuration and Chemical Properties An electron configuration determines its chemical behavior. Outermost shell of these atoms never has more than 4 orbitals or 8 electrons. (2e - /orbital) Electron configuration: Distribution of electrons in an atom s electron shell Electrons must first occupy lower electron shells before the higher shells can be occupied. (A reflection of the natural tendency for matter to move to the lowest possible state of potential energy - the most stable state.) 19 III. Atomic structure determines the behavior of an element (cont.) E. Electron Configuration and Chemical Properties Chemical Properties of an atom depend upon the number of valence electrons. Octet rule: Rule that a valence shell is complete when it contains 8 electrons. (except H & He: outer shell max. = 2 e - ). NOTE: Use this rule only with the smaller elements on the periodic table. Will explain. An atom with a complete valence shell is unreactive or inert. Noble elements (e.g. helium, argon and neon) have filled outer shells in their elemental state and are thus inert. An atom with an incomplete valence shell is chemically reactive (tends to form chemical bonds until it has 8 electrons to fill the valence shell). 20 Valence electrons: Electrons in the outermost energy shell (valence shell)

21 IV. Atoms Combine by Chemical Bonding to form Molecules. Atoms with incomplete valence shells tend to fill those shells by interacting with other atoms. These interactions of electrons among atoms may allow atoms to form chemical bonds. 22 Chemical Bonds: Attractions that hold molecules together. Molecule: (2 + ) atoms held together via chemical bonds. Compound: like a molecule BUT must contain (2 + ) different atoms. A. Covalent Bonds (generally, strongest among the top 3 types of bonds) Covalent bonds are chemical bonds formed by sharing a pair of valence electrons. For example: molecular hydrogen (H 2 ). When 2 hydrogen atoms come close enough for their 1s orbitals to overlap, they share electrons, thus completing the valence shell of each atom.

23 IV. Atoms Combine by Chemical Bonding to form Molecules. A. Covalent Bonds (strongest among the top 3 types of bonds) Structural formula: Formula which represents the atoms and bonding within a molecule (e.g. H-H). The line represents a shared pair of electrons. Molecular formula: Formula which indicates the number and type of atoms (e.g. H 2 ) Single covalent bond: Bond between atoms formed by sharing a single pair of valence electrons. Double covalent bond: Formed when atoms share two pairs of valence electrons (e.g. 0 2 ) 24

25 (SINGLE) 26

IV. Atoms Combine by Chemical Bonding to form Molecules. A. Covalent Bonds (strongest among the top 3 types of bonds) 27 Note that 2 hydrogens are necessary to complete the valence shell of oxygen in water (H 2 O), and four hydrogens are necessary for carbon to complete the valence shell in methane (CH 4 ). Compound: water H 2 O A pure substance composed of two or more (2 + ) elements in a fixed ratio. For example: water (H 2 O) methane (CH 4 ) methane CH 4 28 Why?? DOUBLE Bond actually- 2 double bonds. O = C = O

IV. Atoms Combine by Chemical Bonding to form Molecules. A. Covalent Bonds (strongest among the top 3 types of bonds) 29 Triple covalent bond: Formed when atoms share three pairs of valence electrons (e.g. N 2 ) Test it! ---- Atomic # of N = 7 Valence: Bonding capacity of an atom which is the number of covalent bonds that must be formed to complete the outer shell. Valences of some common elements: Hydrogen = 1 Oxygen = 2 Nitrogen = 3 Carbon = 4 Phosphorous = 3 Sulfur = 2 V. Hydrogen Bonding. Biologically, an important weak bond: Can form between molecules or between different parts of a single large molecule. Example: Integrity of DNA double - stranded molecule; more Hydrogen bond: Bond formed by the charge attraction when a (hydrogen atom covalently bonded to one electronegative (electron-loving) atom) is attracted to (another electronegative atom). 30 Example: NH 3 (ammonia) in H 2 0. Weak attractive force that is about 20 times easier to break than a covalent bond. Is a charge attraction between oppositely charged portions of polar molecules. Can occur between a hydrogen that has a slight positive charge when covalently bonded to an atom with high electronegativity (usually 0 and N).

31 HYDROGEN BOND (dotted line) ELECTRONEGATIVE ATOMS (electron-loving) V. Hydrogen Bonding. 32 Hydrogen Bonding orders water into a higher level of structural organization. + + - Grey = Oxygen atoms + + + + + - + + + - - - The polar molecules of water are held together by hydrogen bonds. Positively charged H of one molecule is attracted to the negatively charged 0 of another water molecule. Each water molecule can form a maximum of four hydrogen bonds with neighboring water molecules.

33 IV. Atoms Combine by Chemical Bonding to form Molecules. B. Ionic Bonds (second strongest among the top 3 types of bonds) Ion: Charged atom or molecule Anion: An atom that has gained one or more electrons from another atom and has become negatively charged; a negatively charged ion. Cation: An atom that has lost one or more electrons and has become positively charged; a positively charged ion. Ionic bond: Bond formed by the electrostatic attraction after the complete transfer of an electron from a donor atom to an acceptor. The acceptor atom attracts the electrons because it is much more electronegative than the donor atom Ionic compounds are called salts. e.g. NaCl or table salt Are strong bonds in crystals, but are fragile bonds in water; salt crystals will readily dissolve in water and dissociate into ions. 34

35 NEXT SLIDE: REVIEW 36 QUICK REVIEW:

VI. Chemical reactions change the composition of matter Chemical reactions: process of making and breaking chemical bonds leading to changes in the composition of matter. The relative concentration of reactants and products affects the reaction rate. (the higher the concentration, the greater probability of a reaction). Process where reactants undergo changes into products. Matter is conserved, so all reactant atoms are only rearranged to form products. Some reactions go to completion (all reactants converted to products), but most reactions are reversible. 37 For example: 3H 2 + N 2 2NH 3 VI. Chemical reactions change the composition of matter Chemical equilibrium: is established when the RATE of the forward reaction equals the RATE of the reverse reaction. 38 Is a dynamic equilibrium with reactions continuing in both directions. Relative concentrations of reactants and products stay the same. NOTE: Chemical equilibrium does NOT mean that the concentrations of reactants and products are equal.

VII. VIII. Water is the solvent of life Solution: A liquid that is a homogenous mixture of two or more substances. Solvent: Dissolving agent of a solution. Solute: Substance dissolved in a solution. Aqueous solution: Solution in which water is the solvent. Organisms are sensitive to changes in ph A. Dissociation of Water Molecules Occasionally, the hydrogen atom that is shared in a hydrogen bond between two water molecules, shifts from the oxygen atom to which it is covalently bonded to the unshared orbitals of the oxygen atom to which it is hydrogen bonded. 39 VIII. Organisms are sensitive to changes in ph A. Dissociation of Water Molecules (cont.) Only a hydrogen ion (proton with a +1 charge) is actually transferred. (H + ) Transferred proton (H + ) binds to an unshared orbital of the second water molecule creating a hydronium ion (H 3 0 + ) 40 Water molecule that lost a proton has a net negative charge and is called an hydroxide ion (OH - ). hydronium ion hydroxide ion H 2 0 + H 2 0 H 3 0 + + OH -

VIII. Organisms are sensitive to changes in ph A. Dissociation of Water Molecules (cont.) 41 By convention, ionization of H 2 0 is expressed as a dissociation into H + and OH -. H 2 0 H + + OH - Reaction is reversible. almost all At equilibrium, most of the H 2 0 is not ionized. VIII. Organisms are sensitive to changes in ph 42 B. Acids and Bases At equilibrium in pure water at 25 o C: Number of H + ions = number of OH - ions. [H + ] = [OH - ] =1/10,000,000 M = 10-7 M Note that brackets indicate molar concentration. Very few water molecules are actually dissociated (only 1 out of 554,000,000 molecules)!!!

ACID Substances that increase the relative [H + ] of a solution. Also removes [OH - ] because it tends to combine with [H + ] to form H 2 O. BASE Substances that reduce the relative [H + ] of a solution. May alternately increase [OH - ] (see NaOH below) For example: A base may reduce [H + ] directly: For example: (in water) HCl [H + ] + Cl - NH 3 + [H + ] NH + 4 A base may reduce [H + ] indirectly: A solution in which: [H + ] = [OH - ] is neutral NaOH [Na + ] + OH - [H + ] > [OH - ] is acidic OH - + H + H 2 O [H + ] < [OH - ] is basic (alkaline) LEO the lion says GER!!! 43 44

VIII. Organisms are sensitive to changes in ph C. The ph Scale A solution in which: [H + ] = [OH - ] is neutral [H + ] > [OH - ] is acidic [H + ] < [OH - ] is basic (alkaline) 45 46

VIII. Organisms are sensitive to changes in ph C. The ph Scale Discuss this on the board ph = -log [H + ] Most biological fluids are within the ph range of 6 to 8. There are some exceptions such as stomach acid with ph = 1.5. Each ph unit represents a tenfold difference (scale is logarithmic), so a slight change in ph represents a large change in actual [H + ]. 47 48 ph Values

VIII. Organisms are sensitive to changes in ph D. Buffers 49 By minimizing wide fluctuations in ph, buffers help organisms maintain the ph of body fluids within the narrow range necessary for life (usually ph 6-8). Buffer: Substance that prevents large, sudden changes in ph. The concept of a buffer will be explained using the following example >>>>>> VIII. Organisms are sensitive to changes in ph D. Buffers Are combinations of H + -donor and H + -acceptor forms of weak acids or bases. Work by accepting H + ions from solution when they are in excess, and by donating H + ions to the solution when they have been depleted. 50 (A buffer in the blood)

IX. Biological Molecules NOTE: Be able to distinguish between organic (contains carbon and hydrogen) and inorganic. I. Most macromolecules (large) are polymers: Polymer: Monomer: Large molecule consisting of many identical or similar subunits connected together. Subunit or building block molecule of a polymer. Macromolecule: (Macro = large) Large organic (carbon containing) polymer. Formation of macromolecules from smaller building block molecules represents another level in the hierarchy of biological organization. 51 IX. Biological Molecules There are four classes of macromolecules in living organisms: 1. Carbohydrates. 2. Lipids. 3. Proteins. 4. Nucleic acids. i.e., DNA, RNA and ATP (which will be reviewed during metabolism). 52 Most polymerization reactions in living organisms are condensation reactions (next slide). Polymerization reactions: Chemical reactions that link two or more small molecules to form larger molecules with repeating structural units.

IX. Biological Molecules Condensation reactions: (ANABOLISM) Polymerization reactions during which monomers are covalently LINKED, producing net removal of a water molecule for each covalent linkage. 53 One monomer loses a hydroxyl (-OH), and the other monomer loses a hydrogen (-H). Process requires energy. Process requires biological catalysts or enzymes. SPLITTING polymers into monomers >>>>> IX. Biological Molecules Hydrolysis: (CATABOLISM) Hydro = water; lysis = break) A reaction process that BREAKS (SPLITS) covalent bonds between monomers by the addition of water molecules. 54 A hydrogen from the water bonds to one monomer, and the hydroxyl (also from the water) bonds to the adjacent monomer. For example: (and many more with illustrations to follow) digestive enzymes catalyze hydrolytic reactions which break apart large food molecules into monomers that can be absorbed into the bloodstream.

55 X. Organisms use CARBOHYDRATES for Fuel and Building Material. CARBOHYDRATES: Organic molecules made of sugars and their polymers. Monomers or building block molecules are simple sugars called monosaccharides. Polymers are formed by condensation reactions. Are classified based upon the number of simple sugars. MONO (1 sugar), DI- (2 sugars) >>>>> X. Organisms use CARBOHYDRATES for Fuel and Building Material. 56 MONOSACCHARIDES: (Mono = single; sacchar- = sugar) Simple sugar in which C, H and 0 occur in the ratio of (CH 2 0) Are major nutrients for cells. Glucose is the most common (a.k.a. as DEXTROSE on an I.V. drip). Store energy in their chemical bonds which is harvested by cellular respiration. Can be incorporated as monomers into disaccharides and polysaccharides.

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X. Organisms use CARBOHYDRATES for Fuel and Building Material. 59 DISACCHARIDES: (DI = two; sacchar- = sugar) A double sugar that consists of two monosaccharides joined by a glycosidic linkage. Glycosidic linkage: Covalent bond formed by a condensation reaction between two sugar monomers. For example, maltose >>>>>>>>>>> X. Organisms use CARBOHYDRATES for Fuel and Building Material. DISACCHARIDES: A condensation!! (removing water to combine---) 60 MALTOSE

X. Organisms use CARBOHYDRATES for Fuel and Building Material. 61 DISACCHARIDES: Some IMPORTANT DISACCHARIDES: 62

63 X. Organisms use CARBOHYDRATES for Fuel and Building Material. POLYSACCHARIDES Macromolecules that are polymers of a few hundred or thousand monosaccharides. Are formed by linking monomers in enzyme-mediated condensation (joining by removing water) reactions. STORAGE Polysaccharides Cells hydrolyze storage polysaccharides into sugars as needed. The most common storage polysaccharide in animals is Glycogen: Glucose polymer in animals. (as starch in plants) Stored in the muscle and liver of humans and other vertebrates. 64

XI. Lipids are mostly hydrophobic molecules with diverse functions LIPIDS: Diverse group of organic compounds that are insoluble in water, but will dissolve in nonpolar solvents (e.g. ether, chloroform, benzene). Important groups are: Fats Phospholipids Steroids A. Fats: Macromolecules constructed from: 1. Glycerol, a three-carbon alcohol. 2. Fatty acid (carboxylic acid). Composed of a carboxyl group (-COOH) head at one end and an attached hydrocarbon chain ("tail"). 65 66

67 Dehydration synthesis H OH +H 2 O H OH +H 2 O H OH +H 2 O 68 XI. Lipids are mostly hydrophobic molecules with diverse functions LIPIDS: A. Fats: Hydrocarbon chain has a long carbon skeleton usually with an even number of carbon atoms (most have 16-18 carbons). Nonpolar C-H bonds make the chain hydrophobic and not water soluble.

What s wrong here? 69 FATTY ACID GLYCEROL ESTER linkage 70 XI. Lipids are mostly hydrophobic molecules with diverse functions LIPIDS: A. Fats: Some characteristics of fat include: Fats are insoluble in water. The source of variation among fat molecules is the fatty acid composition. Fatty acids in a fat may all be the same, or some (or all) may differ. Fatty acids may vary in length. Fatty acids may vary in number and location of carbon - to carbon double bonds (as shown on the next slide) >>>

71 XI. Lipids are mostly hydrophobic molecules with diverse functions LIPIDS: A. Fats: 72

XI. Lipids are mostly hydrophobic molecules with diverse functions LIPIDS: A. Fats: Fat serves many useful functions, such as: Energy storage: One gram of fat stores twice as much energy (calories) as a gram of polysaccharide. (Fat has a higher proportion of energy rich C-H bonds.) More compact fuel reservoir than carbohydrate. Animals store more energy with less weight than plants which use starch, a bulky form of energy storage. Cushions vital organs in mammals (e.g. kidney). Insulates against heat loss (e.g. mammals such as whales and seals). 73 XI. Lipids are mostly hydrophobic molecules with diverse functions LIPIDS: B. Phospholipids: Compounds with molecular building blocks of: Glycerol two fatty acids one phosphate group usually an additional small chemical group attached to the phosphate. 74 Differ from fat in that the third carbon of glycerol is joined to a negatively charged phosphate group.

75 XI. Lipids are mostly hydrophobic molecules with diverse functions LIPIDS: B. Phospholipids: Can have small variable attached to phosphate. Are diverse depending upon differences in fatty acids and in phosphate attachments. Show ambivalent behavior towards H 2 O. Hydrocarbon tails are hydrophobic, and the polar head (phosphate group with attachments) is hydrophilic. Are major constituents of cell membranes. Hydrophilic Head 76 Hydrophobic Tails

XI. Lipids are mostly hydrophobic molecules with diverse functions LIPIDS: C. Steroids: Lipids which have four fused carbon rings with various functional groups attached. Cholesterol, an important steroid: Is the precursor to many other steroids including vertebrate sex hormones and bile acids. Is a common component of animal cell membranes. 77 Can contribute to atherosclerosis. 78 A quick LIPID REVIEW >>>

REVIEW 79 XII. PROTEINS. 80 PROTEINS: A. Overview: A protein is a polymer of amino acids. An amino acid has a carboxyl end (COOH) and an amino end (NH 2 ), as well as a variable R group. Twenty kinds of amino acids are used in protein structure. B. Peptides 1. Two + amino acids is a peptide. 2. A peptide bond is formed between the amino group (NH 2 ) of one amino acid and the carboxyl group (COOH) of the next. 3. Peptide size varies: there are dipeptides, tripeptides, polypeptides.

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XII. PROTEINS. 83 PROTEINS: C. Levels of Protein Structure 1. Primary structure: the order of the amino acids in the peptide. 2. Secondary structure is a coiled or folded shape held together by hydrogen bonds. 3. Tertiary structure is formed by further bending and folding. 4. Quaternary structure: between two or more polypeptide chains. 84

85 Dehydration synthesis H OH +H 2 O A peptide bond (red above) is formed between the amino group (NH 2 ) of one amino acid and the carboxyl group (COOH) of the next. This slide is for illustrative purposes only: 86 Levels of Protein Structure

XII. PROTEINS. PROTEINS: D. Protein Conformation and Denaturation. 1. Protein conformation refers to its overall shape. It cannot function properly if the shape is altered. 2. Denaturing a protein using heat or changes in ph causes it to unravel and often destroys it. E. Protein Functions. Last Slide Protein functions include serving as structural components, for catalysis as ENZYMES, for communication, to provide membrane transport, in cell recognition and protection, and for movement. 87