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New Biology For Engineers and Computer Scientists Introduction & the Chemistry of Life Shu-Ping Lin, Ph.D. Institute of Biomedical Engineering E-mail: [email protected] Website: http://web.nchu.edu.tw/pweb/users/splin/ Course Outline & Evaluation TA: 康毓珊 (BME)e-mail: [email protected] Attendance 10% Homework 10% Quiz 1 10 % Quiz 2 10 % Mid Term 30% Final Exam 30% Office Hours: Thur. 10:00am~ 11:30am (精密大樓3F) Phone:22840734 X22 New Biology for Engineers and Computer Scientists, by Aydin Tozeren and Stephen W. Byers, Publisher: Prentice Hall, 2003. The Chemistry of Life -- Important Elements in Living Organisms Six elements compose 98% of the mass of all living organisms. Hydrogen Carbon Nitrogen Oxygen Phosphorus Sulphur Mineral ions, such as iron, calcium, and magnesium, account for as much as 2 percent of the human body weight. The Chemistry of Life - Atoms The basic unit of each chemical element is the atom. Atoms have a large nucleus, composed of protons and neutrons held together by the Strong Force. The electrons "orbit" the nucleus, attracted by the Electrical Force. The number of protons determines the chemical element, and the number of neutrons determines the isotope of the element. No net electrical charge Ions -- when atoms gain or lose electrons http://cass.ucsd.edu/public/tutorial/scale.html Carbon Orbitals around a nucleus Figure1.1 Periodic Table Periodic table indicating the atomic properties of all elements found on earth. Periodic Table Explorer is a simple Periodic Table software. http://www.technosamrat.com/freewares/periodic-table-explorer/ Chemical Reactivity & Octet Rule Chemical Reactivity – The tendency of a substance to undergo chemical changes in a system The valence electrons are responsible for the combining capacity of atoms. Octet Rule – A tendency for eight electrons in the outermost orbital to undergo chemical changes It is the cardinal rule of bonding. It is the gain in stability when atoms have a full complement of eight electrons in their valence shells. The bonding in carbon dioxide (CO2): all atoms are surrounded by 8 electrons, fulfilling the octet rule http://en.wikipedia.org/wiki/Octet_rule Hydrogen (H) Has single valence electron Form a stable compound hydrogen must either empty its electron orbital by giving it up to another atom (ionic bond) or fill the orbital by sharing electrons with another atom (covalent bond) Form ionic compounds with metals Form molecular compounds with nonmetals Carbon (C) An organic element with the second smallest atomic number Has six protons (Most carbon atoms possess six neutrons.) Has 4 valence electrons (half-filled outermost orbital) allows carbon to form chemical bonds with various other atoms The resulting chemical versatility is essential for reactions of biological metabolism and propagation Nitrogen (N) Has seven protons Has five valence electrons Outer electron orbit only slightly above half filled The compounds of nitrogen, though not as numerous as those of carbon, are just as varied in function. Phosphorus (P) Has fifteen protons Has five valence electrons Abundant in rock formation in the way of calcium phosphate Inorganic salt – Calcium phosphate 60% of bones in the human body Oxygen (O) Has six valence electrons at its outermost shell Has eight protons Abundant in the earth’s crust Gain 2 electrons to achieve the stable octet configuration Sulfur (S) Has six valence electrons Is found in a free-element form in large beds several hundred feet underground Plays a fundamental role in determining the threedimensional shapes of many proteins when present in the amino acid ―cysteine‖ Calcium (Ca) Has 20 valence electrons The most abundant mineral in the human body The average adult body contains in total approximately 1 kg, 99% in the skeleton in the form of calcium phosphate salts. ~22.5 mmol in extracellular fluid (ECF) ~500 mmol of calcium is exchanged between bone and ECF over a period of 24 hours. Ca2+ into and out of the cytoplasm functions as a signal for many cellular processes Covalent Bonds Covalent Bonds – Formed when electron overlaps two atoms by sharing one or more pairs of electrons in order to have eight electrons in the outermost orbitals (octet rule) Only exception – Hydrogen ( has one valence electron) H2: two hydrogen nuclei share the two electrons equally and completely. Each electron is attracted to both protons, but the two protons of the opposing hydrogen atoms repulse each other and the balance of these opposing forces are the determinants of the resulting molecular structure. Sharing of electron pairs in covalent bonds Molecules Molecule – A combination of atoms held together by covalent bonds Identified by the symbols of elements that constitute it Oxygen molecule (O2) – Composed of two atoms Quantity of molecules – One mole per liter (mole/L , or M) The clouds of unshared electrons are larger than those of shared electrons. Charge clouds of nonbonded electrons push the bonded pairs toward each other. Clouds of two unshared pairs are larger than the those of bonded electrons in H2O. H2O: 104.5 ° 3D Structure of an Organic Molecule Shared Pair – The pair of electrons involved in covalent bonding Unshared Pair – The pairs of the outermost electrons not involved in bonding Each bond and unshared pair form a charge cloud that repels all other charge clouds repulsions between the charge clouds determine the 3D shape of the molecules physical rules: Electron pairs spread as far apart as possible to minimize repulsive forces. Since the cloud of unshared pair occupies more space Repulsive Forces between the unshared pairs of electrons > Repulsive Forces H2O: 104.5 ° between two shared pairs Three Commonly Known Molecules Methane, Ammonia, Water Bond Length is not fixed acts as much as a stiff spring Methane the shape of a tetrahedron with bond angles equal to 109.5 ° Ammonia the shape of a tetrahedron with the unshared electron pair occupying one corner of tetrahedron (bond angle: 107°) Water triangular in shape, with the bond angle 104.5° between two hydrogen-oxygen covalent bonds Double & Triple Bonds Single Bonds – Bonds in which a single pair of electrons is shared between two atoms straight-line segment connecting two atoms Double Bonds – Four electrons are shared between two atoms double-line segment Single bond can rotate freely about the bond axis Double bond occupies more space than a single bond more repulsive, Ex: H-C-H bond angle: 116°, H-C=O: 122° Triple Bonds – Sharing of six electrons between two atoms rarely observed in organic materials N-N (unstable) N≡N Organic Compounds – Small Carbon Molecules Double bonds between carbon atoms occur often in biological compounds. Physical Strength of a Covalent Bond Important covalent bonds Covalent bonds constitute stable links between in biological systems: atoms and are the strongest of bonds connecting molecules. Rupture of covalent bonds can occur in two ways: (kcal/mol) Linkage between two atoms are broken symmetrically to provide a pair of free radicals (derived from the unpaired spins of their electrons) often observed in the breaking of identical or similar atoms The bonds are broken asymmetrically to produce a pair of ions electron deficient and electron rich The energy required to break a covalent bond is much greater than the internal (thermal) energy Repulsive forces available at body temperature (0.6 kcal/mol) The energy released by the formation of new bonds between the unshared can break covalent bonds. pairs of electrons make N-N bonds less stable Covalent bonds between carbon atoms are highly than C-C bonds. stable. Electronegativity and Polar Bonds Electronegativity – The capacity of an atom to attract electrons from a neighboring atom measured on a scale from 4 (fluorine, the most electronegative element) to a hypothetical 0 Covalent bonds between two atoms with comparable electronegativity correspond to an equal sharing of electrons between the two nuclei. Polar covalent bond or simply a polar bond – Highly electronegative atoms bond with weaker electronegative atoms partial charges (δ+ and δ-) in different parts of the molecule Unequal sharing of electrons between atoms in a molecule results in charge polarity. Polar molecules dissolve easily in water and interact with other polar molecules. Nonpolar molecules dissolve better in a nonpolar environment, such as acetone or ethanol. Questions Why is sulfur less electronegative than oxygen? Ans: The electro negativity chart shows that oxygen is the most electronegative atom of bioelement, that is to say, oxygen has an electro negativity of 3.5 whereas sulphur has an electro negativity of just 2.5. Therefore, oxygen will take two electrons to fill its s and p sub shells. Even though the charge of four oxygen molecules is -8, each molecule has a -2 charge. What significant role does sulfur play in living organisms? Ans: Sulfur is found in the amino acids methionine, and cysteine. These amino acids are known as the sulfur-bearing amino acids, which are considered the building blocks of all proteins. Sulfur is important for the regulation of plant growth and development since it is the main source used by photosynthetic organisms. Ionic Compounds and Electrostatic Bonds Compound – One of the two interacting atoms is much more electronegative than the other (one or more electrons in the less electronegative atom are transferred to the more electronegative atom) Two electrically charged particles are called ions. Cation – Ion with a positive charge (Ca2+ or H+) Anion – Ion with negative charge (OH-) Ionic Bond – Electrostatic force holds two ions together due to their differing charges. Ionic Compounds – High melting points, conduct electricity in the molten state, and tend to be soluble in water (Na+Cl- and Ca32+(PO43-)2) Ammonia (NH3) – Nitrogen forms three covalent bonds with 3 hydrogen atoms one of the outermost electron pairs of nitrogen is not shared dissolved in water, ammonia picks up a hydrogen ion (H+) shared a previously unshared bond to become an ammonium ion (NH4+), has a net positive charge of 1 Ethanol – Has no charge, but ionized ethanol has a charge of -1 due to acidity of the medium Water and Hydrogen Bonds At RT, oxygen and hydrogen: gaseous form; water molecule: liquid state due to Hydrogen Bonding between water molecules Water – Has important physical properties that make life possible on the Earth, ability to dissolve many other substances, serves as a medium in which a great variety of chemical changes occur Hydrogen Bonding – Caused by the polar nature of covalent H—O bonds that hold together water molecules Hydrogen atoms of water molecules are attracted to the unshared electrons of oxygen atoms of adjacent water molecules. Hydrogen bonds (3-dots): Water molecules form transient hydrogen bonds with several others, creating a fluid network. Hydrogen Bonds Water molecules form hydrogen bonds with other polar molecules including those of nitrogen and oxygen compounds. Nitrogen and oxygen are more electronegative than hydrogen Covalent N—H and O—H bonds are polar bonds, the H atoms in these bonds can participate in hydrogen bonding. Amino (—NH2) and hydroxyl (—OH) groups Two atomic groups that often engage in hydrogen bonding in living systems The presence of these two groups makes many molecules soluble in water Water molecules cluster around cations (Na+) and anions (Cl-) in solutions, blocking their association into a solid. Hydrogen Bonds – Also form between different parts of large polar molecules such as proteins, much weaker than a covalent bond in physical strength and duration, greatly influence the physical properties of biological substances Hydrogen bonds between methanol (CH3OH) and water Hydrogen bonds between methylamine (CH3NH2) and water Van der Waals Attraction Hydrophobic or water fearing – Inability of nonpolar hydrocarbon (compounds of hydrogen and carbon, nonpolar because of sharing bonding electrons equally) molecules to form hydrogen bonds with water Hydrophobic effect – Drives a number of very important biological phenomen – Formation of cell membranes by lipid bilayers Hydrophobic nature of the hydrocarbon chains of the lipids make up the bilayer Lipid bilayers have a hydrophobic core with the hydrophilic heads facing the polar environments outside or inside the cell Drives the folding of proteins Hydrophobic amino acids (building blocks of proteins) tend to fold away from more polar regions Form a hydrophobic pockets or clefts Provide a binding site for small hydrophobic molecules such as steroid hormones Van der Waals interactions – Brief and weak attraction, Van der Waals forces act on nonpolar molecules brought together by a polar solvent. Nonpolar molecules disturb the hydrogen bonds between water molecules, forcing water to form a cage around them Hydrophobic tails extend into the air to avoid contact with water Types of Bonds between Molecules, Atoms or Ions *Four types of bonds between atoms or ions: Ionic Bonding – Electrons are transferred from one atom to another producing a charge on both atoms (ions), in order to keep them together. NaCl Covalent Bonding – Sharing of electrons between atoms resulting in an overlap of their electron orbitals. Diamond Metallic Bonding – Formed when closely packed atoms bond by sharing electrons from their inner electron shells, whereas the outer shell electrons are free to form a sea of electrons Giving metals their high electrical and thermal conductivities Van der Waals Bonds – Weak electrostatic interactions that hold adjacent sheets of talc or graphite together, act on nonpolar molecules brought together by a polar solvent. *Bonds between molecules: Hydrogen bonding – Caused by the polar nature of covalent H—O bonds that hold together water molecules Acids and Bases Concentration of hydrogen ions in a solution: pH= -log[H+], pH of cell fluid (cytoplasm)= 7.2~ 7.3 Acid – Donate protons (H+), pH<7, sour taste, and cause the purple dye litmus to turn red, Ex: HCl H+ and Cl-, —COOH (carboxyl group molecules) H+ and COO-, gastric fluids, vinegar, soda, lemon juice, and contracting muscle cell fluid Base – Accept protons, pH>7, taste bitter, feel slippery, and turn the purple dye litmus blue, Ex: NH3 NH4+, —NH2 (amino group in biological molecules) —NH3+, sweat, human blood ([H+]=4×10-8M), and household ammonia Functional groups in biological molecules undergo acid-base reaction in living systems water H+ and OH-; H+ affects the rates of chemical reactions Cellular pH play an important role in cell division and cell growth. Many proteins denature in acidic conditions except gastric enzymes and lysosomal enzymes. Cellular membranes pump protons from one side to the other Create a pH gradient An important component in the generation and storage of energy in mitochondria Chemical Reactions Occurs when atoms combine or change binding partners Arrow in the following equation The direction of the chemical reaction The number of atoms on the left-hand side of the equation = The number of atoms on the right-hand side C3 H8 5O2 3CO2 4H 2O Reactants Products Energetics of Chemical Reactions In thermodynamics – ―System‖: the part of the universe that is of interest, ―surroundings‖: the rest of the universe Open system – Exchange matter and energy with its surrounding, Ex: living cells (take up nutrients, release waste products, and generate work and heat 1st law of thermodynamics: the conservation of energy law and states that energy can be neither destroyed nor created Energy gained by an open system U q w Work done by the system on the surroundings Heat absorbed by the surroundings (release heat q<0: Exothermic process Endothermic) 2nd law of thermodynamics: some of the energy involved irreversibly loses its ability to do work; S is entropy Entropies of the system S1 S2 Su 0 Entropies of the surroundings Spontaneous reaction Entropies of the universe Chemical Reactions Combine the 1st and 2nd laws of thermodynamics H G TS △H: Overall change in bond energy due to reaction (kcal/mol) △S: Change in entropy (measure of the energy lost to disorder in system) △G: Change in free energy Free energy (△G) – Energy can be used to do work For spontaneous chemical processes occur at constant temperature S q / T q: Heat imported into the system 1 T: Absolute temperature Spontaneity for constant pressure and temperature processes, the criterion of △G < 0 A + B C + D △G < 0, exergonic or energetically favorable, △G products<△G reactants , spontaneously forward reaction △G > 0, endergonic, reverse reaction △G = 0, chemical equilibrium (at steady state), both forward and reverse reactions occur at equal rates Free Energy Change (△G) in a Chemical Reaction 2 parts to the free-energy change that occurs in a chemical reaction – △G0 (Change in standard free energy), Concentrations of reactants and products G G o RT ln{([C ][ D]) /([ A][ B])} R(gas constant)=1.978 cal/mol, T=273.15 °+ ℃ △G increases with increasing temperature and increasing concentrations of reactants Chemical Reactions in Biological Systems Many chemical reactions in biological systems are endergonic reactions △G>0 reverse reaction How to proceed in the forward direction? Additive property of free energy Chemical reaction with positive △G may be coupled to a reaction with a negative △G of a larger magnitude A B + D △Go1 =12 kcal/mol D X + Y △Go2 =-16 kcal/mol △G3= △Go1+ △Go2=- 4 kcal/mol <0 exergonic (reactions operate in the forward direction) This exergonic reaction coupled the hydrolysis of ATP ATP: The Standard Energy Carrier Nucleotide monophosphate Nucleoside Pi Pi H++PO43- ATP: nucleotide adenine, sugar ribose, and triphosphate unit Participates in most cellular reactions and processes that require energy Very stable molecule ATP hydrolysis reaction by the enzyme – ATPase, releases 7.2 kcal/mol ATP + H2O ADP+ Pi + H+ PO43-