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

CS 790 – Bioinformatics. A Gentle Introduction to (or review of) Fundamentals of Chemistry and Organic Chemistry. Square one…. Isotopes of Chlorine Atomic Natural Isotope Protons Neutrons mass abundance 35 Cl 17 18 34.97 76% 37 Cl 17 20 36.97 24%. Fundamentals of Chemistry.

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

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  1. CS 790 – Bioinformatics A Gentle Introduction to(or review of)Fundamentals of Chemistryand Organic Chemistry Square one…

  2. Isotopes of Chlorine Atomic Natural Isotope Protons Neutrons mass abundance 35Cl 17 18 34.97 76% 37Cl 17 20 36.97 24% Fundamentals of Chemistry • Reading the periodic table • Neutrons and isotopes • 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 6CCarbon12.01 Intro to biochemistry

  3. Types of Orbitals Second Letter Number Maximum quantum denoting of number of number orbitals orbitals electrons 0 s 1 2 1 p 3 6 2 d 5 10 3 f 7 14 Electron shells and orbitals • Quantum numbers • n = First quantum number = shell • l = Second quantum number = orbital type • Golden rule: l < n Know these two. Intro to biochemistry

  4. # electrons in the subshell 2p5 Electron shell Type of orbitals Electron Subshells 1st Quantum 2nd Quantum Notation fornumber number subshells1 0 1s 2 0,1 2s,2p3 0,1,2 3s,3p,3d4 0,1,2,3 4s,4p,4d,4f … Subshells and valence • All orbitals of the same type (same l and n) are called a subshell • Subshellnotation: Intro to biochemistry

  5. Electronic configurations • 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 Intro to biochemistry

  6. 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 Intro to biochemistry

  7. F F F Covalent bonds: Lewis diagrams • How many covalent bonds will an atom form? • Flourine: Atomic number = 9, Electron configuration: 1s2,2s2,2p5 • Oxygen: Atomic number = 8 Electron configuration: 1s2,2s2,2p4 F F or O O O O O or Intro to biochemistry

  8. 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. Intro to biochemistry

  9. Ions and ionic bonds • Formation of ions • Conflicting goals: neutral charge vs. stable electronic configuration • Some atoms have a strong tendency to gain or lose electrons: • Sodium (Na): Atomic # = 11: 1s2,2s2,2p6,3s1  Na+ • Chlorine (Cl): A# = 17: 1s2,2s2,2p6,3s2 ,3p5  Cl– • Complete electron transfer, no sharing • Coulombs law: • Ionic bond or salt bridge Intro to biochemistry

  10. 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+ I – Intro to biochemistry

  11. 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 andcovalent bonds • Donor: covalent/polar bond to H • Acceptor: ionic attraction to H H+ O – H+ O N Intro to biochemistry

  12. Van der Waals bonds • Nonspecific – when any two atoms at ~3 to 4 Å apart • Å = angstrom units = 1010 meters = 0.1 nm • Low energy interaction • Significantly smaller thanh-bonds or ionic attraction • Adds up over many atoms • When two atoms have very similar shapes, the Van derWaals contacts can become significant Intro to biochemistry

  13. 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 Intro to biochemistry

  14. Looking at chemical structures Propane: Benzene: H H H H H C C C H C C H C C H H H H H C C CH3 CH2 CH3 H H C C C Intro to biochemistry

  15. CH2 CH2 CH3 CH3 CH CH3 CH3 CH3 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 C4H10 • Butane: • Isobutane: Intro to biochemistry

  16. 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 Intro to biochemistry

  17. Some Common Functional Groups Intro to biochemistry

  18. Concentration • 1 mole of a substance = 6.02 × 1023 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 Na2SO4) and add enough water to make 1 liter of solution: M = [Na2SO4] = 1.0 Intro to biochemistry

  19. 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 H3O+ • Strong acids dissociate completely • Weak acids do not dissociate completely • pH of a solution • pH = log[H+] Intro to biochemistry

  20. More on pH • A simple example: • Suppose we add 0.001 moles of HCl to 1.0 L of H20 • [H+] = 103 moles/liter, so pH = 3 • 0 7 14acidic basic • Bases accept H+ ions • pOH = log[OH ] • pH + pOH = 14 • Water: pH = 7, pOH = 7 Intro to biochemistry

  21. 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 Intro to biochemistry

  22. Properties of Water • The polarity of water makes it highly cohesive: • Water solvates & weakensionic and hydrogen bonds: Intro to biochemistry

  23. Hydrophobic Attraction • Nonpolar (hydrophobic atoms), are driven together • Hydrophobic interactions • Driven by water’s affinity for itself Intro to biochemistry

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