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Chapter 1: Structure and Bonding

Chapter 1: Structure and Bonding. Organic Chemistry – Mrs. Meer 2011-2012 www.mtlsd.org/teachers/smeer. vancomycin C 66 H 75 Cl 2 N 9 O 24. vancomycin C 66 H 75 Cl 2 N 9 O 24. Compounds you may know. Chapter 1 Objectives.

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Chapter 1: Structure and Bonding

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  1. Chapter 1: Structure and Bonding Organic Chemistry – Mrs. Meer 2011-2012 www.mtlsd.org/teachers/smeer

  2. vancomycin C66H75Cl2N9O24

  3. vancomycin C66H75Cl2N9O24

  4. Compounds you may know

  5. Chapter 1 Objectives • Review material from first year chemistry such as atomic structure and chemical bonding • Determine the correct Lewis structure for basic organic molecules using VSEPR theory • Understand hybrid orbitals and determine which ones would be present in each molecule • Interconvert between line-angle and Lewis structures

  6. History of Organic Chemistry Organic Chemistry – the chemistry of carbon compounds Swedish chemist Torbern Bergman (in 1770) was the first to distinguish organic compounds (compounds coming from living things) from inorganic compounds. Organic compounds were believed to have a “vital force”. Michel Chevreul (in 1816) showed that organic compounds, without a vital force, could be turned into other organic compounds when making soap. Friedrich Wöhler (in 1828) isolated an “organic” compound from “inorganic” material, accidentally! No More Vitalism!!

  7. Usefulness of Organic Chemistry Why is organic useful to you? You are organic (DNA, proteins, carbohydrates, lipids, etc.) Biological sciences (medicine, pharmacy, etc.) Polymer chemistry Food chemistry Nanotechnology …survive college organic and do well on the MCATs

  8. What makes carbon unique? There are more carbon compounds than compounds of all the other elements combined. Carbon forms very strong (high bond energy) covalent bonds with many different types of atoms allowing long chains to form. Carbon can be found as many “allotropes” or coaxed into different arrangements (Note: Not all compounds listed below are considered organic.) Rings (cyclo- compounds) Graphite (pure C) Diamond (pure C) Buckeyballs (spheres of pure C) Nanotubes (tubes of pure C) Polymers (Styrofoam®)

  9. …so many possibilities… ethanol dimethyl ether These are isomers.

  10. Open the Odyssey Program on the student computers • Open the ‘Molecular Labs’ section • Under the ‘Organic’ section, open Lab 58, “The Bonding Characteristics of Carbon” • Complete the worksheet that goes along with it.

  11. Atomic Structure Structure of the atom: made of protons, neutrons and electrons atomic number (Z) – number of p+ mass number (M) – number of p+ + no example: 12C example: hydrogen-2 atoms are neutral, so p+ = e-

  12. Types of Species Isotopes – atoms with the same number of protons yet different masses or mass numbers (different number of neutrons) Ex. 12C, 13C, 14C for carbon and 1H, 2H, 3H for hydrogen Ions – atoms with the same number of protons yet different charges (different number of electrons) increases stability of atoms (octet rule) Ex. Na+, O2-, Fe3+

  13. Metals vs. Nonmetals • Metals form cations. Na Na+ + 1e- • Nonmetals form anions. Cl + 1 e- Cl-

  14. Charge determination (WITH SOME EXCEPTIONS!) • Group 1 – forms 1+ • Group 2 – forms 2+ • Group 13(B and Al) – forms 3+ • Group 15 – forms 3- • Group 16 – forms 2- • Group 17 – forms 1- • Group 18 – doesn’t forms ions easily!

  15. Noble gases are very stable and don’t react. Every element on the periodic table will try to react to be stable, like the noble gases.

  16. Electronic Structure of the Atom An element’s reactivity is dependent upon its electrons - electrons take part in bonding. Electrons show particle-wave duality. paddle-wheel experiment (particle nature) double-slit experiment (wave nature)

  17. Electronic Structure of the Atom Electrons are found in orbitals. An orbital is a region of space where there is a high probability of finding an electron. It is designated by ψ2(ψ is a wave with a + and – sign, so we use ψ2 so we’re always dealing in the positive region).

  18. Orbitals – the 1s and 2s Recall quantum numbers: The first two energy levels can hold 2 and 8 electrons, respectively First level: 1s-orbital Electron density is a function of distance from the nucleus. Highest density is at the nucleus. Second level: 2s-orbital, 2p-orbital Region of space where there is no electron density is a node. Most of the density is farther away, so 2s is higher in energy than the 1s.

  19. 1s, 2s, 3s

  20. p orbitals p orbitals: px, py, pz 6 electrons total 3 orientations (all degenerate ) p orbitals are in the 2nd, 3rd, 4th, 5th, and 6th energy levels

  21. p orbitals px, py and pz are called degenerate orbitals because they have equal energy. p orbitals are higher in energy because average electron density is farther than the 1s or 2s. p orbitals are for n = 2-7

  22. Other orbitals d, f, and g orbitals exist, but we don’t worry about them in organic chemistry since we deal with carbon. carbon – 1s22s22p2

  23. Pauli Exclusion Principle – a maximum of 2 e- may occupy one orbital, both with opposites spins Aufbau Principle – fill in the lowest possible energy orbital Aufbau Principle – fill in the lowest possible energy orbital Orbital Diagrams Hund’s Rule – within equal energy orbitals, the e- are distributed to have the maxiumum unpaired e- possible 3d 4s Pauli Exclusion Principle – a maximum of 2 e- may occupy one orbital, both with opposite spins Aufbau Principle – fill in the lowest possible energy orbital 3p 3s 2p 2s Energy increases as you go up. 1s

  24. degenerate orbitals (same E)

  25. Section 1.3 Electronic Configurations

  26. Section 1.3 Electronic Configurations

  27. Section 1.3 Electronic Configurations What is the electron configuration for: carbon – • oxygen – • manganese – • lead –

  28. Section 1.3 Electronic Configurations What is the electron configuration for: carbon – 1s22s22p2 • oxygen – 1s22s22p4 • manganese – 1s22s22p63s23p64s23d5 • lead – 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2

  29. Valence Electrons valence electrons - electrons in the highest energy level Br: 1s22s22p63s23p64s23d104p5 1st E level 2nd E level 3rd E level 4th E level 4th is the highest Energy level, so there are 7 valence electrons

  30. You Try It How many valence electrons do the following have? Mg 2 C 4 S 6 F 7 H 1 N 5 O 6

  31. Development of Chemical Bonding Theory August Kekulé and Archibald Couper proposed that carbon is tetravalent (1858). Carbon always forms 4 bonds to make a stable compound. Multiple bonding was proposed when Emil Erlenmeyer showed acetylene (C2H2) to have a triple bond (1862) and Crum Brown showed ethylene (C2H4) to have a double bond (1864). August Kekulé determined that carbon chains can link end to end to become rings, such as benzene (1865).

  32. Development of Chemical Bonding Theory Jacobus van’t Hoff proposed that the four bonds of carbon are not random, but are three-dimensionally arranged with specific direction (1874). He helped determine that the hydrogens in methane are positioned at the corners of a tetrahedron, 109.5o.

  33. Bond Formation: The Octet Rule G.N. Lewis (1916) proposed theories about how atoms form bonds Atoms transfer or share electrons in such a way as to attain a filled valence shell of electrons (the Octet Rule). Covalent bonding involves the sharing of electrons. Equal sharing: non-polar bond; Ex: C-C or C-H Unequal sharing: polar bond; Ex: C-O or O-H Ionic bonding involves the loss of an electron due to a large difference in electronegativity (>2.0).

  34. The two fundamental types of bonds. Pure Covalent Ionic

  35. There is another type of bond, not purely covalent and not purely ionic. Pure Covalent Nonpolar Covalent Polar Covalent Ionic

  36. Electronegativity • electronegativity – a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound The difference in electronegativity values for two atoms will indicate whether the two atoms form an ionic bond or a polar or nonpolar covalent bond. What is the general trend for electronegativity?

  37. Electronegativity

  38. Covalent Bonding Elements that have similar electronegativities will share electrons. • We draw the sharing of two electrons with a line – single bond • We draw the sharing of four electrons with a double line = double bond • We draw the sharing of six electrons with a triple line = triple bond

  39. Valence Shell Electron Pair Repulsion Theory (VSEPR Theory) • Where do the electrons come from to make the bond? • They are valence electrons. • Usually one valence electron comes from each atom to form a covalent bond. • octet rule – Atoms transfer or share electrons in such a way as to attain a filled valence shell of electrons (8). • What would the compound C2H6 look like?

  40. Region of Electron Density • Regions of electron density spread out three-dimensionally around an atom due to repulsions. • Types of regions of electron density • single bond • double bond • triple bond • lone pair

  41. General Rules for Drawing Lewis Structures • Add up total number of valence electrons (the ones that bond) • Determine the central atom • Determine bonding scheme (HONC-1234) • Distribute remaining electrons to follow the octet rule (lone pairs)

  42. HONC - 1234

  43. Lewis Structures – Single Bonding In a Lewis structure, each valence electron is symbolized by a dot, bonding electrons are symbolized by lines (non-bonding electrons are drawn as a pair of dots). Try to arrange ALL of the following compounds so they have a noble gas configuration (full octet). Draw the Lewis structures for CH4 CH3NH2 CH3CH2OH CH3Cl

  44. Lewis Structures – Multiple Bonds The sharing of one pair of electrons is called a single bond. Sharing of 2 pairs is a double bond. Sharing of 3 pairs is a triple bond. In some compounds, the ONLY way to satisfy all element’s octets is to use multiple bonds. Ex. Draw Lewis structures for C2H4, CH2O and C2H2

  45. Lewis Structures – Single Bonding • Try these: • Ex. CH3CH2COCH2NH2 • Ex. CH3CO2CH=CHCH3 • Ex. CH3CHOHCH2CONHCH2CH3

  46. Lewis Structures – More Complex Ex. CH3(CH2)3COOH Ex. CH3CHOHCH2CONHCH2CH3 Ex. CH3CO2C(CH3)=CHCH3

  47. Valence Bond Theory Valence Bond Theory – a covalent bond forms when two atoms approach each other closely and a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom. Why? An overlap of the s and s orbitals is called a s bond. How many s bonds are in H2?

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