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Chapter 1 Chemical Bonding and Chemical Structure

Chapter 1 Chemical Bonding and Chemical Structure. The branch of chemistry that deals with carbon based compounds Organic compounds may contain any number of other elements, including hydrogen, nitrogen, oxygen, halogens, phosphorus, silicon, and sulfur. Organic chemistry. Methane. Sucrose.

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Chapter 1 Chemical Bonding and Chemical Structure

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  1. Chapter 1Chemical Bonding and Chemical Structure

  2. The branch of chemistry that deals with carbon based compounds Organic compounds may contain any number of other elements, including hydrogen, nitrogen, oxygen, halogens, phosphorus, silicon, and sulfur Organic chemistry Methane Sucrose Morphine

  3. History • Vitalism: Only biological systems (e.g., plants, animals) could produce organic compounds • Wohler’s synthesis of urea (1828), began to undermine vitalism

  4. Why Study Organic Chemistry? • Organic chemistry lies at the heart of the modern chemical industry • Central to medicine and pharmacy • Interface of physical and biological sciences • Everyday applications: Plastics, textiles, communications, transportation, food, clothing, cosmetics, etc.

  5. Review of Chemical Bonding • Valence Electrons: Outermost electrons • s and p electrons for main group elements • Responsible for chemical properties of atoms • Participate in chemical reactions Valence Electron Core Electrons

  6. Octet Rule • Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells • Natural electron configuration of the Noble Gases • Done by gaining, losing, or sharing electrons • Increases stability • H and He seek a “Duet”

  7. Ionic Bonding • Ions: atoms that have a charge due to gain or loss of electrons • Anion: (-) charged atom • Cation: (+) charged atom • Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

  8. Formula Unit

  9. Ionic bonds are omni-directional • Can dissociate into free ions

  10. Covalent Compounds • Covalent Compounds: compounds composed of atoms bonded to each other through the sharing of electrons • Electrons NOT transferred • No + or – charges on atoms • Non-metal + Non-metal • Also called “molecules” • Examples: • H2O • CO2 • Cl2 • CH4

  11. orH-H Duet or

  12. Covalent Bonds

  13. Electronegativity • The measure of the ability of an atom to attract electrons to itself • Increases across period (left to right) and • Decreases down group (top to bottom) • fluorine is the most electronegative element • francium is the least electronegative element

  14. Electronegativity Scale

  15. ENCl = 3.0 3.0 - 3.0 = 0 Pure Covalent Types of Bonding • Non-Polar Covalent Bond: • Difference in electronegativity values of atoms is 0.0 – 0.4 • Electrons in molecule are equally shared • Examples: Cl2, H2, CH4

  16. ENCl = 3.0 ENH = 2.1 3.0 – 2.1 = 0.9 Polar Covalent • Polar Covalent Bond: • Difference in electronegativity values of atoms is 0.4 – 1.7/2.0 • Electrons in the molecule are not equally shared • The atom with the higher EN value pulls the electron cloud towards itself • Partial charges • Examples: HCl, ClF, NO

  17. Electrostatic Potential Maps • A graphical depiction of electron distribution

  18. ENCl = 3.0 ENNa = 1.0 3.0 – 0.9 = 2.1 Ionic • Ionic Bond: • Difference in EN above 1.7-2.0 • Complete transfer of electron(s) • Whole charges

  19. Dipole Moment (m) • Depends on charge separation and distance • m = qr (a vector quantity) • q = magnitude of charge • r = vector from site of + charge to site of – charge • Units = Debyes (D)

  20. Molecular Polarity

  21. Lewis Dot Structures • Count the number of valence electrons present in the molecule • Determine the arrangement of atoms. Generally, the atom that occurs least often is central. Join the terminal atoms to the central atom(s) using shared pairs of electrons (bonds) • Place any remaining electrons around the terminal atoms to satisfy the octet rule • Exception: Hydrogen • Place any remaining electrons on the central atom(s) to satisfy the octet rule

  22. Check to make sure: • You’ve used the correct number of valence electrons • Everyone has an octet (or duet) • Everyone is doing what they like to do • If the number of electrons around the central atom is less than 3, change the single bonds to multiple bonds

  23. What Things Like To Do • Halogens • Like to be terminal • Like to have one bonding pair (two shared electrons) and 3 lone pairs (non-bonding electrons) • Carbon • Likes to have 4 bonding pairs and no lone pairs • Likes to bond to other carbons • Likes to be central • Silicon • Likes to do what carbon does • Notice, it sits under C on the periodic table

  24. Oxygen • Like to have 2 bonding pairs and 2 lone pairs • Sulfur • Likes to do what O does • Nitrogen • Likes to have 3 bonding pairs and 1 lone pair • Phosphorous • Likes to do what N does

  25. Hydrogen • Likes to be terminal with only 1 bond • Do not put lone pairs on H • Boron • Likes to have 3 bonds and no lone pairs • Likes a sextet instead of an octet (what everybody else besides Hydrogen likes) • *Note: • A double bond = 2 bonding pairs • A triple bond = 3 bonding pairs

  26. Problems • Draw the Lewis Dot Structures for the following molecules • CO2 • P2H4 • O3 • NO3-

  27. Drawing Resonance Structures -1 -1 • Draw first Lewis structure that maximizes octets • Assign formal charges • Move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge

  28. Formal Charge • Assigned charge for each atom in a molecule/ion • Electronic bookkeeping – may or may not correspond to a real charge • Sum of formal charges on each atom must equal the total charge on the molecule/ion • FC = Valence e-’s – Lone Pair e-’s – ½ bonding e-’s

  29. Molecular Structures of Covalent Compounds • Atomic connectivity:How atoms in a molecule are connected • Molecular geometry: How far apart atoms are and how they are arranged in space • Bond lengths • Bond angles • Dihedral angles OR

  30. Bond Length • Distance between nuclei

  31. Increases with atoms in higher rows • Decreases toward higher atomic number along a row • Decreases with increasing bond order

  32. Bond Angles • Angle between each pair of bonds • Contribute to molecular shape • Determined by Valence-shell electron-pair repulsion (VSEPR) • Use molecular models! • Line-and-wedge structures

  33. Drawing LDS With Correct Geometry

  34. Valence Shell Electron Pair Repulsion Theory • VSEPR theory: • Electrons repel each other • Electrons arrange in a molecule themselves so as to be as far apart as possible • Minimize repulsion • Determines molecular geometry

  35. Defining Molecular Shape • Electron pair geometry: the geometrical arrangement of electron groups around a central atom • Look at all bonding and non-bonding e-’s • Molecular Geometry: the geometrical arrangement of atoms around a central atom • Ignore lone pair electrons

  36. Problems • Predict the approximate geometry in each of the following molecules • BF3 • HCN • CO32- • Estimate the bond angles and relative bond lengths in the following molecule

  37. Dihedral Angle Also known as the torsional angle Rotation can occur along single bonds 48

  38. Valence Bond Theory

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