A Gentle Introduction to (or Review of ) Fundamentals of Chemistry and Organic Chemistry

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Wright State University CORE Scholar Computer Science and Engineering Faculty Publications Computer Science and Engineering 2003 A Gentle Introduction to (or Review of ) Fundamentals of Chemistry and

1 Wright State University CORE Scholar Computer Science and Engineering Faculty Publications Computer Science and Engineering 2003 A Gentle Introduction to (or Review of ) Fundamentals of Chemistry and Organic Chemistry Dan E. Krane Wright State University - Main Campus, dan.krane@wright.edu Michael L. Raymer Wright State University - Main Campus, michael.raymer@wright.edu Follow this and additional works at: Part of the Computer Sciences Commons, and the Engineering Commons Repository Citation Krane, D. E., & Raymer, M. L. (2003). A Gentle Introduction to (or Review of) Fundamentals of Chemistry and Organic Chemistry.. This Presentation is brought to you for free and open access by Wright State University s CORE Scholar. It has been accepted for inclusion in Computer Science and Engineering Faculty Publications by an authorized administrator of CORE Scholar. For more information, please contact corescholar@

2 CS 790 Bioinformatics A Gentle Introduction to (or review of) Fundamentals of Chemistry and Organic Chemistry Square one CS790 Bioinformatics

3 Fundamentals of Chemistry Reading the periodic table Neutrons and isotopes Isotopes of Chlorine Atomic Natural Isotope Protons Neutrons mass abundance 35 Cl % 37 Cl % Electron shells, subshells and orbitals Each orbital can hold at most 2 electrons In the ground state orbitals are filled from lower to higher energy 6 C Carbon

4 Electron shells and orbitals Quantum numbers n = First quantum number = shell l = Second quantum number = orbital type Golden rule: l < n Types of Orbitals Second Letter Number Maximum quantum denoting of number of number orbitals orbitals electrons 0 s p d f 7 14 Know these two. 3

5 Subshells and valence All orbitals of the same type (same l and n) are called a subshell Subshell notation: Electron shell 2p 5 Electron Subshells 1 st Quantum 2 nd Quantum Notation for number number subshells 1 0 1s 2 0,1 2s,2p 3 0,1,2 3s,3p,3d 4 0,1,2,3 4s,4p,4d,4f # electrons in the subshell Type of orbitals 4

6 Since the subshells are filled from lowest to highest energy, we can specify only the outermost shell. Atoms tend to lose or gain electrons such that the outermost subshell is full: valence Electronic configurations 5

7 Covalent Bonds For almost all of the elements that we will deal with, 8 valence electrons is an electronically stable configuration. Covalent bonds are formed when atoms share electrons to fill the valence shell 6

8 Covalent bonds: Lewis diagrams How many covalent bonds will an atom form? Flourine: Atomic number = 9, Electron configuration: 1s 2,2s 2,2p 5 F F F or F F Oxygen: Atomic number = 8 Electron configuration: 1s 2,2s 2,2p 4 O O O or O O 7

9 How many covalent bonds? Note the common valences for the elements most common in proteins and DNA: Carbon Oxygen Nitrogen Hydrogen Sulfur Note the similarity between S and O. 8

10 Formation of ions Ions and ionic bonds Conflicting goals: neutral charge vs. stable electronic configuration Some atoms have a strong tendency to gain or lose electrons: Sodium (Na): Atomic # = 11: 1s 2,2s 2,2p 6,3s 1 Na + Chlorine (Cl): A# = 17: 1s 2,2s 2,2p 6,3s 2,3p 5 Cl Complete electron transfer, no sharing q q force = + 2 d Coulombs law: Ionic bond or salt bridge 9

11 Polar Bonds In reality, some atoms will attract shared electrons more strongly. That is, the shared electrons will be off center. The tendency to attract electrons is called electronegativity. There is a continuum between covalent bonds and ionic bonds. K I K δ+ I δ 10

12 The Hydrogen Bond When hydrogen forms a polar bond, the nucleus is left without any unshared electrons It can make a secondary bond with another negative ion, called a hydrogen bond Very common in water: Weaker than polar and covalent bonds Donor: covalent/polar bond to H Acceptor: ionic attraction to H H δ+ δ O H δ+ O N 11

13 Van der Waals bonds Nonspecific when any two atoms at ~3 to 4 Å apart Å = angstrom units = meters = 0.1 nm Low energy interaction Significantly smaller than h-bonds or ionic attraction Adds up over many atoms When two atoms have very similar shapes, the Van der Waals contacts can become significant 12

14 Energy of molecular interactions 1 calorie = the amount of energy to raise the temperature of 1g of water from 14.5 to 15.5 C Molecules have about 0.6 kcal/mole of energy from heat/vibration Molecular interactions: C C : 83 kcal/mole Electrostatic and hydrogen bonds: ~3 7 kcal/mole Van der Walls interaction: ~1 kcal/mole 13

15 Looking at chemical structures Propane: Benzene: H H H H C C C H H H H H H C C C C C H C H CH 3 CH 2 CH 3 H H C C C 14

16 A hydrocarbon isomer Carbon can make 4 covalent bonds There are more carbon-based compounds present on earth than the total of all compounds lacking carbon We could spend an entire course examining the properties of hydrocarbons: molecules made up only of carbon and hydrogen. Example: Isomers of C 4 H 10 Butane: CH 3 CH 2 CH 2 CH 3 Isobutane: CH 3 CH CH 3 CH 3 15

17 Double Bonds Double bonds can force a molecule or functional group to be planar: Geometric isomers cis = on the same side trans = on the opposite side 16

18 Some Common Functional Groups 17

19 Concentration 1 mole of a substance = atoms or molecules of that substance C atomic weight = 12, one mole = 12 grams We express concentration in molarity or moles/liter. Denoted [x]. Example If we take 1 mole of sodium sulfate (142.1g of Na 2 SO 4 ) and add enough water to make 1 liter of solution: M = [Na 2 SO 4 ] =

20 Acids and Bases Acids give off protons in solution HCl H + + Cl In water, the H + ion often binds with water to form a hydronium ion H 3 O + Strong acids dissociate completely Weak acids do not dissociate completely ph of a solution ph = log[h + ] 19

A simple example: More on ph Suppose we add 0.001 moles of HCl to 1.

21 A simple example: More on ph Suppose we add moles of HCl to 1.0 L of H 2 0 [H + ] = 10 3 moles/liter, so ph = acidic basic Bases accept H + ions poh = log[oh ] ph + poh = 14 Water: ph = 7, poh = 7 20

22 pka For a weak acid, the pka is a measure of the tendency of the acid to dissociate (give of an H + ion) Key rule: ph = pka : protonated and unprotonated forms are at equilibrium ph < pka : more protonated ph > pka : less protonated Biological ph varies but is generally close to neutral (7.0) or slightly acidic 21

23 Properties of Water The polarity of water makes it highly cohesive: Water solvates & weakens ionic and hydrogen bonds: 22

24 Hydrophobic Attraction Nonpolar (hydrophobic atoms), are driven together Hydrophobic interactions Driven by water s affinity for itself 23

may contain one or more neutrons

may contain one or more neutrons Biology 115 Fall 2001 Campos/Saupe Atoms and Molecules I. Introduction - living things are composed of the same chemical elements as the nonliving world and obey the same physical and chemical laws - living