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Chapter 12. Saturated Hydrocarbons

Chapter 12. Saturated Hydrocarbons. Sections.

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Chapter 12. Saturated Hydrocarbons

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  1. Chapter 12. Saturated Hydrocarbons Sections

  2. 12.1 Organic and Inorganic Compounds12.2 Bonding Characteristics of the Carbon Atom12.3 Hydrocarbons and Hydrocarbon Derivatives12.4 Alkanes: Acyclic Saturated Hydrocarbons12.5 Structural Formulas12.6 Alkane Isomerism12.7 Conformations of Alkanes12.8 IUPAC Nomenclature for Alkanes12.9 Line-Angle Formulas for Alkanes12.10 Classification of Carbon Atoms12.11 Branched-Chain Alkyl Groups12.12 Cycloalkanes12.13 IUPAC Nomenclature for Cycloalkanes12.14 Isomerism in Cycloalkanes12.15 Sources of Alkanes and Cycloalkanes12.16 Physical Properties of Alkanes and Cycloalkanes12.17 Chemical Properties of Alkanes and CycloalkanesChemistry at a Glance: Properties of Alkanes and Cycloalkanes

  3. Introduction • Organic chemistry is the study of the compounds of carbon atoms consisting of single, double, and triple covalent bonds with C, H, O, N, and some metals • Biochemistry is the study of the compounds of carbon atoms consisting of single, double, and triple covalent bonds with C, H, O, N, and some metals

  4. 12.1 Organic and Inorganic Compounds

  5. Sections • Chapter 1: Covalent Bonding and Shapes of Molecules • Introduction • Electronic Structure of atoms • Lewis Model of bonding • Bond angles and molecular shapes • Polar and non-polar molecules • Resonance • Orbital overlap of covalent bonding • Functional groups

  6. Introduction • Organic chemistry is the study of the compounds of carbon. • C is a small atom • it forms four bonds consisting of single, double, and triple bonds • it forms strong bonds with C, H, O, N, and some metals • Problem: Identify organic and inorganic compounds: • NaCl, C2H5OH, CH3COOH, Na2CO3, CH3OK, KOH

  7. Atomic Orbitals • s orbital - a spherical-shaped atomic orbital; can hold a maximum of 2 electrons • p orbital - a dumbbell-shaped atomic orbital; the three p orbitals (px, py, pz) can hold a maximum of 2 electrons each • Electrons always fill starting with the lowest-energy orbital: • lower energy higher energy • 1s2 2s2 2p6 3s2 3p6 • We will be concerned with only the valence electrons which are the outermost electrons involved in forming bonds.

  8. Shapes of s Atomic Orbitals • All s orbitals have the shape of a sphere, with its center at the nucleus • of the s orbitals, a 1s orbital is the smallest, a 2s orbital is larger, and a 3s orbital is larger still

  9. Shapes of p Atomic Orbitals • A p orbital consists of two lobes arranged in a straight line with the center at the nucleus

  10. Electronic Structure of atoms (Review) • Ground state electronic configuration of atoms in core format • Carbon (C): ): [He] 2s2, 2p2 • or[He] 2s2, 2px13py13pz0 • Potassium (K): {Ar] 4s1 • Phosphorous (P): [Ne] 3s2, 3p3 • Valence shell electronic configuration • Carbon (C): ): 3s2, 3p2 • Potassium (K): 4s1 • Phosphorous (P): 3s2, 3p3 • How you get the electronic configuration of an atom from the periodic table?

  11. Excited State Valence Electron Configuration • Carbon (C): • Ground state: 2s2, 2p2 or2s2, 2px13py13pz0 • Excited State: 2s1, 2p3 or 2s1, 2px13py13pz1

  12. Lewis Model of bonding (Review) • "octet rule“ • atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell • “Lewis structure of atoms” • Shows only valence electrons, is a convenient way of representing atoms to show their chemical bonding pattern.

  13. Lewis structure of atoms (Review)

  14. Cations and Anions • Cations • What elements lose electrons? And how many? • What is the positive charge on their cations? • Anions • What elements gain electrons? • What is the positive charge on their anions? • Covalent bonds • How many covalent bonds are formed? • What elements share electrons?

  15. Covalent or Ionic • Identify covalent and ionic compounds: • NaCl, C2H5OH, CH3COOH, Na2CO3, CH3OK, KOH • Covalent : • Ionic:

  16. Ionic model of bonding model(Review) • Ionic bond - results from the electrostatic attraction between a cation and an anion of two atoms typically involves a metal and a nonmetallic element. • Anion: An atom that gains electrons becomes a negative ion • Cation: An atom that loses electrons becomes a positive ion

  17. Covalent model of bonding(Review) • Covalent bonds - results from the sharing of electrons between two atoms typically involves two nonmetallic elements

  18. Electronegativity (Review) • H • 2.1 • Li Be B C N O F • 1.0 1.5 2.0 2.5 3.0 3.5 4.0 • Na Mg Al Si P S Cl • 0.9 1.2 1.5 1.8 2.1 2.5 3.0 Is the attraction of an atom for its valence electrons increases i n c r e a s e s

  19. Nonpolar/polar-covalent and ionic bonds. (Review) • We classify chemical bonds as polar covalent, nonpolar covalent and ionic based on the difference in electronegativity between the atoms

  20. Classify following bonds nonpolar-covalent, polar-covalent or ionic bonds • N-H nonpolar-covalent, polar-covalent or ionic bonds • O-H nonpolar-covalent, polar-covalent or ionic bonds • C-H nonpolar-covalent, polar-covalent or ionic bonds • C-F nonpolar-covalent, polar-covalent or ionic bonds • Na-Cl nonpolar-covalent, polar-covalent or ionic bonds • Al-Cl nonpolar-covalent, polar-covalent or ionic bonds

  21. Drawing Lewis structure molecules and ions(Review) • 1)Predict arrangement of atoms. • H is always a terminal atom. • Halogens and oxygen are oftenterminal. • The central atom of binary compounds is usually written • first and has the lowest subscript. • Most organic compounds have more than two central atoms. • These are mainly C, but N, O and S can also be central atoms. 2) Total number of valence electrons (e-) • Add all valence electron of atoms in the molecule from the formula. • Add the ion charge for negative ions or subtract for positive ions.

  22. Drawing Lewis structure molecules and ions (Review) • 3) Draw the skeletal structure by connecting the atoms with single bonds. 4) Give each of the atoms an octet (8 e-). Adding unshared pairs of electrons 5) Count the total number of e- used through step 4 and compare to the number calculated in step 2. • a) If it results in zero, the structure is correct. • b) For every two electrons too many, another bond is added (minimize formal charges). • Multiple bonds form only with C, N, O and S. • Total number of bonds to neutral atoms: 4 bonds to C 3 bonds to N, P 2 bonds to O, S 1 bond to H, F, Cl, Br, I

  23. Calculation of formal charges of atoms in the Lewis structure • 1. For a neutral molecule, the sum of the formal charges equals zero. For a polyatomic ion, the sum of the formal charges equals the charge on the ion. • 2. Formal charge of each atom is calculated by: • (group #) - (# unshared e-) - ½ (# shared e-) • 3. Formal charges are shown as + or - on the atom with that charge. • 4. An atom with the same number of bonds as its group number has no formal charge. • 5. In a molecule if two different elements can be assigned a negative charge, then the more electronegative element gets the charge; the same sign should not be given to bonded atoms.

  24. Types of electrons • Bonding pairs • Two electrons that are shared between two atoms. • A covalent bond. • Unshared (nonbonding ) pairs • A pair of electrons that are not shared between two atoms. Lone pairs or nonbonding electrons. Unshared pair oo H Cl oo oo oo Bonding pair

  25. Lewis Structures(Review) • How many bonding electron pairs are in the molecule? • How many bonding electron pairs are in each atom? • How many nonbonding electron pairs are in the molecule?

  26. Draw Lewis structure of molecules • CHCl3 • C2H4 • C3H8O • CH3CH2CH2OH • CH3CH2OCH3 • CH3CO2H • CH3CHO

  27. Draw Lewis structure and assign formal charges • CH3NH3+ • CH3O-

  28. Valence-Shell Electron-Pair Repulsion (VSEPR) model(Review) • For predicting shapes of molecules and polyatomic Ions based on the repulsion of valence pairs of electrons making them as far apart as possible around an atom of a Lewis structure. • 1) Draw the Lewis structure for the molecule or ion. 2) Determine the number of bonding and unshared pairs attached to the central atom. One single, double or triple bond counted as a bonding pair • 3) Choose the appropriate case from the given chart.

  29. Predict the bond angles of molecules from their Lewis structures. (Review)

  30. Polar and nonpolar molecules

  31. Molecular Shape and Polarity (Review)

  32. Curved arrow Electron pushing • Curved arrow:a symbol used to show the redistribution of valence electrons • In using curved arrows, there are only two allowed types of electron redistribution: • from a bond to an adjacent atom • from an atom to an adjacent bond • Electron pushing by the use of curved arrows is also used in explaining reaction mechanisms

  33. Drawing Curved Arrows To show the movement of electrons in breaking and forming bonds. The tail of the arrow is started at the site of electron density (negative character such as a pi bond or lone pair of electrons) and proceeds to the arrowhead which is drawn to the site of electron deficiency (positive character). NEGATIVE TO POSITIVE! Arrows can be drawn from: 1) lone pair bond lone pair 2) bond bond 3) bond

  34. Resonance For many molecules and ions with double bonds, two or more Lewis structure could be written

  35. Orbital Overlap Model

  36. Hybrid Atomic Orbitals Hybridization is the mixing up of two or more atomic orbitals • There are three types of hybrid atomic orbitals for carbon sp3(one s orbital + three p orbitals give four sp3 orbitals) sp2(one s orbital + two p orbitals give three sp2 orbitals) sp (one s orbital + one p orbital give two sp orbitals)

  37. s and p hybrids Four sp3 hybrids Three sp2 hybrids Two sp hybrids

  38.  and  bonds • Overlap of hybrid orbitals can form two types of bonds, depending on the geometry of the overlap • bondsare formed by “direct” overlap  bondsare formed by “parallel” overlap of unhybrid p prbitlas

  39.  and  bonds in single and multiple bonds • single bond - one shared pair of electrons between two atoms; a s bond • double bond - two shared pairs of electrons between two atoms; one s bond and one p bond • triple bond - three shared pairs of electrons between two atoms; one s bond and two p bonds

  40. Bond Properties • Bond strength: • strongest weakest • Bond length: • longest shortest

  41. Counting  and  bonds in Lewis structure

  42. Predicting hybridization of atoms in a Lewis structure • Count sigma bonds and unshared electrons around the atom If the total number of pairs: 2 sp hybridization 3 sp2 hybridization 4 sp3 hybridization

  43. Functional Groups in Organic Compounds • Functional group: an atom or group of atoms within a molecule that shows a characteristic set of physical and chemical properties • Functional group • divide organic compounds into classes • the sites of characteristic chemical reactions • the basis for naming organic compounds

  44. Common Functional Groups

  45. Common Functional Groups (continued)

  46. Classification of organic compounds Class Functional group Example

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