Chapter 1 Structures and Bonding - PDF

John E. McMurry www.cengage.com/chemistry/mcmurry Chapter 1 Structure and Bonding Edited by: Dr. Ahmed Khalil Department of Chemistry, College of Science King Faisal University, Saudi Arabia Email: [email protected]  What is...

John E. McMurry www.cengage.com/chemistry/mcmurry Chapter 1 Structure and Bonding Edited by: Dr. Ahmed Khalil Department of Chemistry, College of Science King Faisal University, Saudi Arabia Email: [email protected]  What is Organic Chemistry? § Living things are made of organic chemicals (carbon-based compounds) § Proteins that make up hair § DNA, controls genetic make-up § Foods, medicines  Origins of Organic Chemistry Foundations of organic chemistry from mid-1700’s. Compounds obtained from plants, animals hard to isolate, and purify. Compounds also decomposed more easily. Torben Bergman (1770) first to make distinction between organic and inorganic chemistry. It was thought that organic compounds must contain some “vital force” because they were from living sources.  Origins of Organic Chemistry Because of “vital force”, it was thought that organic compounds could not be synthesized in laboratory like inorganic compounds. 1816, Chevreul showed that not to be the case, he could prepare soap from animal fat and an alkali and glycerol is a product 1828, Woehler showed that it was possible to convert inorganic salt ammonium cyanate into organic substance “urea”  Origins of Organic Chemistry § Organic chemistry is study of carbon compounds. § Why is it so special? § 90% of more than 30 million chemical compounds contain carbon. § Examination of carbon in periodic chart answers some of these questions. § Carbon is group 4A element, it can share 4 valence electrons and form 4 covalent bonds.  Abundance of Organic Compounds Why are there so many more organic compounds than inorganic? Carbon has unique bonding characteristics – Strong, covalent bonds with C and H Isomerism – Groups of carbon atoms can form more than one unique compound Structural Isomers of C2H6O H H H H H C C O H H C O C H H H H H  Important Topics in General Chemistry 1 Students should revise the following parts in General Chemistry 1 course: § Atomic Structure § Atomic Number and Atomic Mass § Shells, Subshells, and Orbitals § Atomic Structures: Orbitals, Shapes § Electron Configuration: Aufbau Principle, Pauli exclusion principle, and Hund's rule Bonding Characteristics of Carbon C Q: If 2s electrons are already paired, with only 2 2p electrons Valence 2p unpaired, how shell electrons does carbon form 2s 4 covalent bonds? 1s  Development of Chemical Bonding Theory § Kekulé and Couper independently observed that carbon always has four bonds § van't Hoff and Le Bel proposed that the four bonds of carbon have specific spatial directions § Atoms surround carbon as corners of a tetrahedron van't Hoff and Le Bel  Development of Chemical Bonding Theory § Atoms form bonds because the compound that results is more stable than the separate atoms § Ionic bonds in salts form as a result of electron transfers § Organic compounds have covalent bonds from sharing electrons (G. N. Lewis, 1916)  Development of Chemical Bonding Theory § Lewis structures (electron dot) show valence electrons of an atom as dots § Hydrogen has one dot, representing its 1s electron § Carbon has four dots (2s2 2p2) § Kekulé structures (line-bond structures) have a line drawn between two atoms indicating a 2 electrons covalent bond. § Stable molecule results at completed shell, octet (eight dots) for main-group atoms (two for hydrogen)  Development of Chemical Bonding Theory § Atoms with one, two, or three valence electrons form one, two, or three bonds. § Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet. § Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4).  Development of Chemical Bonding Theory § Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3). § Oxygen has six valence electrons (2s2 2p4) but forms two bonds (H2O) Development of Chemical Bonding Theory  Non-Bonding Electrons § Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons § Nitrogen atom in ammonia (NH3) § Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair  Describing Chemical Bonds: Valence Bond Theory § Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom § Two models to describe covalent bonding. Valence bond theory, Molecular orbital theory Valence Bond Theory: § Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms § H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals § H-H bond is cylindrically symmetrical, sigma (s) bond  Bond Energy § Reaction 2 H· ® H2 releases 436 kJ/mol § Product has 436 kJ/mol less energy than two atoms: H–H has bond strength of 436 kJ/mol. (1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)  Bond Energy § Distance between nuclei that leads to maximum stability § If too close, they repel because both are positively charged § If too far apart, bonding is weak  Hybridization of Atomic Orbitals. Pure atomic 4 valence e-1 orbitals. C..C... Lewis symbol.C. Hybrid orbitals. The mathematical process of replacing pure atomic orbitals with reformulated atomic orbitals for bonded atoms is called hybridization, and the new orbitals are called hybrid orbitals.  Hybridization: sp3 Orbitals and the Structure of Methane § Carbon has 4 valence electrons (2s2 2p2) § In CH4, all C–H bonds are identical (tetrahedral) § sp3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931) McMurry Organic Chemistry 6th edition Chapter 1 20 (c) 2003  Tetrahedral Structure of Methane § sp3 orbitals on C overlap with 1s orbitals on 4 H atom to form four identical C-H bonds § Each C–H bond has a strength of 438 kJ/mol and length of 110 pm § Bond angle: each H–C–H is 109.5°, the tetrahedral angle. McMurry Organic Chemistry 6th edition Chapter 1 21 (c) 2003  Hybridization: sp3 Orbitals and the Structure of Ethane § Two C’s bond to each other by s overlap of an sp3 orbital from each § Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds § C–H bond strength in ethane 420 kJ/mol § C–C bond is 154 pm long and strength is 376 kJ/mol § All bond angles of ethane are tetrahedral McMurry Organic Chemistry 6th edition Chapter 1 22 (c) 2003  Hybridization: sp2 Orbitals and the Structure of Ethylene § sp2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp2) § sp2 orbitals are in a plane with120°angles § Remaining p orbital is perpendicular to the plane Ethylene 90° 120° McMurry Organic Chemistry 6th edition Chapter 1 23 (c) 2003  Bonds From sp 2 Hybrid Orbitals § Two sp2-hybridized orbitals overlap to form a s bond § p orbitals overlap side-to-side to formation a pi (p) bond § sp2–sp2 s bond and 2p–2p p bond result in sharing four electrons and formation of C-C double bond § Electrons in the s bond are centered between nuclei § Electrons in the p bond occupy regions are on either side of a line between nuclei McMurry Organic Chemistry 6th edition Chapter 1 24 (c) 2003  Structure of Ethylene § H atoms form s bonds with four sp2 orbitals § H–C–H and H–C–C bond angles of about 120° § C–C double bond in ethylene shorter and stronger than single bond in ethane § Ethylene C=C bond length 133 pm (C–C 154 pm) McMurry Organic Chemistry 6th edition Chapter 1 25 (c) 2003  Hybridization: sp Orbitals and the Structure of Acetylene § C-C a triple bond sharing six electrons § Carbon 2s orbital hybridizes with a single p orbital giving two sp hybrids § two p orbitals remain unchanged § sp orbitals are linear, 180°apart on x-axis § Two p orbitals are perpendicular on the y-axis and the z-axis McMurry Organic Chemistry 6th edition Chapter 1 26 (c) 2003  Orbitals of Acetylene § Two sp hybrid orbitals from each C form sp–sp s bond § pz orbitals from each C form a pz–pz p bond by sideways overlap and py orbitals overlap similarly McMurry Organic Chemistry 6th edition Chapter 1 27 (c) 2003  Bonding in Acetylene § Sharing of six electrons forms CºC § Two sp orbitals form s bonds with hydrogens McMurry Organic Chemistry 6th edition Chapter 1 28 (c) 2003  Hybridization of Nitrogen and Oxygen § Elements other than C can have hybridized orbitals § H–N–H bond angle in ammonia (NH3) 107.3° § N’s orbitals (sppp) hybridize to form four sp3 orbitals § One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H McMurry Organic Chemistry 6th edition Chapter 1 29 (c) 2003  Hybridization of Oxygen in Water § The oxygen atom is sp3-hybridized § Oxygen has six valence-shell electrons but forms only two covalent bonds, leaving two lone pairs § The H–O–H bond angle is 104.5° McMurry Organic Chemistry 6th edition Chapter 1 30 (c) 2003  Molecular Orbital Theory § A molecular orbital (MO): where electrons are most likely to be found (specific energy and general shape) in a molecule § Additive combination (bonding) MO is lower in energy § Subtractive combination (antibonding) forms MO is higher McMurry Organic Chemistry 6th edition Chapter 1 32 (c) 2003  Molecular Orbitals in Ethylene § The p bonding MO is from combining p orbital lobes with the same algebraic sign § The p antibonding MO is from combining lobes with opposite signs § Only bonding MO is occupied McMurry Organic Chemistry 6th edition Chapter 1 33 (c) 2003  Representing Organic Compounds § Compounds can be represented in many different ways. § Some representations provide information about the structure, while others do not. § Molecular Formula: Number of atoms of each element in one molecule of a compound (no structural information). § Ex. C6H14 § Empirical Formula: Relative ratio’s of elements present. § Ex. CH2O could be CH2O or C2H4O2. § Line bond (Kekule) Structures: Show all atoms and bonds. § Condensed Structures: Show all atoms, but only show bonds when necessary. § Skeletal Structures: Show C-C bonds and all atoms that are not C or H. Representing Organic Compounds H H H H Structural H C C C C H Also called Line-bond Formula or Kekule structures. H H H H Condensed CH3CH2CH2CH3 Structural Formula CH3(CH2)2 CH3 Skeletal Learn this quickly, it Structure will save you much time. Molecular C4H10 Formula  Summary § Organic chemistry – chemistry of carbon compounds § Covalent bonds - electron pair is shared between atoms § Valence bond theory - electron sharing occurs by overlap of two atomic orbitals § Molecular orbital (MO) theory, - bonds result from combination of atomic orbitals to give molecular orbitals, which belong to the entire molecule § Sigma (s) bonds - Circular cross-section and are formed by head-on interaction § Pi (p) bonds – “dumbbell” shape from sideways interaction of p orbitals § Carbon uses hybrid orbitals to form bonds in organic molecules. § In single bonds with tetrahedral geometry, carbon has four sp3 hybrid orbitals § In double bonds with planar geometry, carbon uses three equivalent sp2 hybrid orbitals and one unhybridized p orbital § Carbon uses two equivalent sp hybrid orbitals to form a triple bond with linear geometry, with two unhybridized p orbitals § Atoms such as nitrogen and oxygen hybridize to form strong, oriented bonds § The nitrogen atom in ammonia and the oxygen atom in water are sp383- hybridized