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Chapter 1 Introduction • Organic Chemistry - is the study of carbon compounds (ranging from simple compounds such as CH4 to complex ones such as proteins, carbohydrates, DNA, etc…). • Torbern Bergman (1770) – was the first to make a distinction between organic compounds (“those that can only be derived from living things”) and inorganic compounds (“those that were found in the nonliving world”). Vitalism The belief in a life force outside the physical and chemical laws. • Michel-Eugene Chevreul (1816) – converted an organic substance (fat) to other organic substances (fatty acids and glycerin) without the help of an outside vital force. Animal fat NaOH Soap + Glycerin H2O Soap H3O+ “Fatty acids” • Friedrich Wohler (1828) – accidentally converted the inorganic salt ammonium cyanate into the organic substance urea. • Stanley Miller (1953) – used electrical discharges to trigger reactions in a primitive “atmosphere” of H2O, H2, NH3 (ammonia), and CH4 (methane). His experiment demonstrated that the spontaneous synthesis of organic compounds could have been an early stage in the origin of life. Mechanism The belief that all natural processes are governed by physical and chemical laws. • Vitalism - is the belief in a life force outside the physical and chemical laws. • Mechanism - is the belief that all natural processes are governed by physical and chemical laws. Carbon Atoms: The Most Versatile Building Blocks of Molecules • Carbon is tetravalent. It has 4 electrons in its valence shell. • Carbon completes its valence shell by sharing electrons with other atoms in four covalent bonds. Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings • Carbon can form a covalent bond with another carbon atom, forming the carbon skeleton of organic compounds (e.g. C2H6, C2H4, …). Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings • Carbon can bond to a variety of atoms, including O, H, and N (e.g. CO2, urea CO(NH2)2…) Carbon dioxide, CO2 O C O O H Urea, CO(NH2)2 C N N H H H • Variation in carbon skeletons contributes to the diversity of organic molecules. • Carbon skeletons vary in: Length Shape and arrangement (branching or ring) Presence and location of double and/or triple bonds Fig. 4.4. Variations in carbon skeletons Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings • Hydrocarbons - are organic molecules consisting only of carbon (C) and hydrogen (H) atoms. Hydrocarbons Aliphatic Hydrocarbons Alkanes Alkenes Aromatic Hydrocarbons Alkynes I. Atomic Structure A. Overview B. Orbitals C. Electron Configuration A. Overview Atoms are composed of smaller parts, called subatomic particles: Protons (p+) Neutrons (n0) Electrons (e-) Protons (p+) are positively charged are tightly packed with the neutrons to form a dense core, called the atomic nucleus (10-1410-15 m in diameter), at the center of the atom have a mass of 1 dalton (amu = 1.7 x 10-24 g) Neutrons (n0) are uncharged are tightly packed with the protons to form the atomic nucleus at the center of the atom have a mass of 1 dalton (amu = 1.7 x 10-24 g) Electrons (e-) are negatively charged move around the nucleus at a distance of about 10-10 m and at a speed of light forming a cloud have a mass of 1/2000 dalton Structure of an atom Positively charged nucleus (protons and neutrons) is very dense and small (10-15 m) Negatively charged electrons form a cloud (10-10 m) around nucleus • Diameter is about 2 10-10 m (200 picometers (pm)) [1 Å (angstrom) = 10-10 m = 100 pm] Atomic Number and Mass Number An atom is described by its: • Atomic number (Z) = # of protons • Mass number (A) = # of protons + # of neutrons All the atoms of a given element have the same atomic number. Isotopes - are atoms of the same element with different numbers of neutrons and thus different mass numbers. Atomic Mass - is the weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes (= atomic weight). B. Orbitals Orbital is a three-dimensional space where an electron is found 90-95% of the time. Each electron orbital holds up two electrons. How are electrons distributed in an atom? • Erwin Schrödinger (1926) – proposed a wave equation, a mathematical expression based on the concept that electrons show properties not only of particles but also of waves. Quantum Mechanics A branch of science that studies particles and their associated waves. • Quantum mechanics: describes electron energies and locations by a wave equation. • The solution to Schrödinger equation led to a wave function, y. • Wave function, y - gives the shape of the electron orbital. • Each wave function,y, is an orbital. • The square of the wave function y2 - gives the probability of finding the electron. y2 gives the electron density for the atom. y2 is called the probability density. • Electron cloud has no specific boundary so the most probable area is shown. Shapes of Atomic Orbitals for Electrons • Based on those derived for a hydrogen atom, there are four different kinds of orbitals for electrons: s, p, d, and f. • s and p orbitals are most important in organic chemistry. spherical, nucleus at center dumbbell-shaped, nucleus at middle cloverleaf-shaped, nucleus at middle Orbitals and Shells • Orbitals are grouped in shells of increasing size and energy • Different shells contain different numbers and kinds of orbitals • Each orbital can be occupied by two electrons 3rd shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons 2nd shell contains one s orbital (2s) and three p orbitals (2p), eight electrons 1st shell contains one s orbital, 1s, and holds only two electrons p-Orbitals • In each shell there are three perpendicular p orbitals, px, py, and pz, of equal energy. • Lobes of a p orbital are separated by region of zero electron density, a node. C. Electron Configuration Ground-state electron configuration of an atom: • is a list of orbitals occupied by its electrons • is the lowest-energy arrangement • can be determined by following three “rules” Rules for Electron Configuration 1. The Aufbau Principle. Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d 2. The Pauli Exclusion Principle. No more than two electrons can occupy an orbital, and they must be of opposite spin, and . 3. Hund’s Rule. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron. Practice Problem: Give the ground-state electron configuration for each of the following elements: a. Boron b. Phosphorus c. Oxygen d. Chlorine Practice Problem: How many electrons does each of the following elements have in its outermost electron shell? a. Potassium b. Aluminum c. Krypton II. Chemical Bonding A. Development of Chemical Bonding Theory B. Nature of Chemical Bonding A. Development of Chemical Bonding Theory • Kekulé and Couper (1858) independently observed that carbon always has four bonds. • Extensions to the Kekulé-Couper theory include: • Multiple bonding • Ring formation • van't Hoff and Le Bel (1874) proposed that: The four bonds of carbon have specific spatial directions. Atoms surround carbon as corners of a tetrahedron. Note that a dashed line indicates a bond is behind the page Note that a wedge indicates a bond is coming forward Practice Problem: Draw a molecule of chloroform, CHCl3, using solid, wedged, and dashed lines to show its tetrahedral geometry Practice Problem: Convert the following representation of ethane, C2H6, into a conventional drawing that uses solid, wedged, and dashed lines to indicate tetrahedral geometry around each carbon B. Nature of Chemical Bonding • Atoms form bonds because the compound that results is more stable (has less energy) than the separate atoms. • Stable molecule results at completed shell, octet (eight dots) for main-group atoms (two for hydrogen). Octet Rule • Atoms often gain, lose, or share electrons to achieve a configuration of the noble gas closest to them. • Atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). Major types of Bonds 1. Ionic Bond - results from total electron transfer. - is an electrostatic attraction between ions. 2. Covalent Bond - results from electron sharing. - first described by G. N. Lewis (1916). Lewis Structures (or electron-dot structures) Valence electrons – are those that reside in the incomplete, outermost orbital of an atom (valence shell). Lewis structures - show valence electrons of an atom as dots. The # of electrons available for bonding are indicated by unpaired dots. Drawing Lewis Structures 1. Add the valence electrons from all atoms. 2. Identify the central atom (usually the one with the highest molecular mass and closest to the center of the periodic table). 3. Place the central atom in the center of the molecule and add all other atoms around it. 4. Place one bond (two electrons) between each pair of bonded atoms. 5. Complete the octet for all atoms surrounding the central atom. 6. Complete the octet for the central atom. Use double or triple bonds if necessary. Number of Covalent Bonds to an Atom • 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. Valences of Carbon • Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4). Valences of Oxygen • Oxygen has six valence electrons (2s2 2p4) but forms two bonds (H2O). Valences of Nitrogen • Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3). Non-bonding electrons • Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons. Example: Nitrogen atom in ammonia (NH3) Kekulé Structures (or line-bond structures) Kekulé structures - show a two-electron covalent bond as a line drawn between atoms. Lonepairs of nonbonding valence electrons are often not shown. Practice Problem: What are likely formulas for the following substances? a. GeCl? b. AlH? c. CH?Cl2 d. SiF? e. CH3NH? Practice Problem: Write both Lewis and line-bond structures for the following substances, showing all nonbonding electrons: a. CHCl3, chloroform b. H2S, hydrogen sulfide c. CH3NH2, methylamine d. NaH, sodium hydride e. CH3Li, methyl lithium Practice Problem: Why can’t an organic molecule have the formula C2H7? III. Covalent Bonding A. Valence Bond Theory B. Hybridization C. Molecular Orbital Theory A. Valence Bond Theory Valence Bond Theory is one of two models that describes covalent bond formation. describes a covalent bond as resulting from the overlap of two atomic orbitals 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. • Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms. Valence Bond Theory • 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 2H: H–H has bond strength of 436 kJ/mol. (1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ) Bond Length • It is the 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 B. Hybridization • Hybridization - is a process through which a hybrid orbital is formed. • Hybrid orbital - is an orbital derived from a combination of atomic orbitals. • Hybrid orbitals, such as sp3, sp2, and sp hybrids of C, are strongly directed and form stronger bonds than unhybridized atomic orbitals, s or p. Hybridization: sp3 Orbitals and the Structure of Methane • Carbon has 4 valence electrons (2s2 2p2) • In CH4, all C–H bonds are identical (tetrahedral) even though C uses two kinds of orbitals. How does C form four identical C–H bonds in CH4? • sp3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Linus Pauling (1931) Why does C form four identical C–H bonds in CH4? • When s orbital hybridizes with three p orbitals, the resultant sp3 hybrid orbitals are unsymmetrical about the nucleus and strongly oriented. • This is due to opposite algebraic signs of p lobes. • This allows to form strong bonds. Tetrahedral Structure of Methane • sp3 orbitals on C overlap with 1s orbitals on 4 H atoms 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. 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 three H’s 1s orbitals to form six C–H bonds Hybridization: sp3 Orbitals and the Structure of Ethane • 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 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 with 120° angles • Remaining p orbital is perpendicular to the plane 90 120 Bonds from sp2 Hybrid Orbitals • Two sp2-hybridized orbitals overlap to form a sigma (s) bond • p orbitals overlap side-to-side to form a pi () bond Bonds from sp2 Hybrid Orbitals • sp2–sp2 s bond and 2p–2p bond share four electrons and form C-C double bond Electrons in the s bond are centered between nuclei Electrons in the bond occupy regions on either side of a line between nuclei Structure of Ethylene • H atoms form s bonds with four sp2 orbitals • H–C–H and H–C–C bond angles are about 120° • C=C double bond in ethylene is shorter and stronger than single bond in ethane • Ethylene C=C bond length 133 pm (Ethane C–C 154 pm) Hybridization: sp Orbitals and the Structure of Acetylene • sp hybrid orbitals: Carbon 2s orbital hybridizes with a single p orbital giving two sp hybrids • two p orbitals remain unchanged Hybridization: sp Orbitals and the Structure of Acetylene • sp orbitals are linear, 180° apart on x-axis • Two p orbitals are perpendicular on the y-axis and the z-axis Orbitals of Acetylene • Two sp hybrid orbitals from each C form sp–sp s bond • pz orbitals from each C form a pz–pz bond by sideways overlap and py orbitals overlap similarly to form py–py bond Bonding in Acetylene • Sharing of six electrons forms C C • Two sp orbitals form s bonds with hydrogens Hybridization: Nitrogen and Oxygen • Atoms such as nitrogen and oxygen hybridize to form strong, oriented bonds • Nitrogen in NH3 • Oxygen in H2O Hybridization of Nitrogen in Ammonia • H–N–H bond angle in ammonia (NH3) is 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 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° To assign hybridization: • Draw the Lewis structure for the molecule or ion • Determine the electron domain geometry • Specify the hybrid orbitals needed to accommodate the electron pairs based on their geometric arrangement – Example: the hybrid orbitals used by N in NH3 molecule are predicted by first writing the Lewis structure, and finally the geometry. To assign hybridization: # attached atoms + # nonbonding electron pairs (nondelocalized) Hybridization Bond Angles Molecular Geometry 2 sp 180o Linear 3 sp2 120o Trigonal Planar 109.5o Tetrahedral, pyramidal, or bent 4 sp3 Practice Problem: Draw a line-bond structure for propane, CH3CH2CH3. Predict the value of each bond angle, and indicate the overall shape of the molecule Practice Problem: Convert the following molecular model of hexane, a component of gasoline, into a linebond structure Practice Problem: Draw a line-bond structure for propene, CH3CH=CH2, indicate the hybridization of each carbon, and predict the value of each bond angle Practice Problem: Draw a line-bond structure for 1,3-butadiene H2C=CH-CH=CH2: indicate the hybridization each carbon, and predict the value of each bond angle Practice Problem: Draw both a Lewis structure and a line-bond structure for acetaldehyde, CH3CHO Practice Problem: Shown below is a molecular model of aspirin (acetylsalicyclic acid). Identify the hybridization of each carbon atom, and tell which atoms have lone pairs of electrons Practice Problem: Draw a line-bond structure for propyne, CH3CCH2, indicate the hybridization of each carbon, and predict the value of each bond angle Practice Problem: Draw Lewis and line-bond structures for formaldimine,CH2NH. How many electrons are shared in the carbon-nitrogen bond? What is the hybridization of the nitrogen atom? Practice Problem: What geometry do you expect for each of the following atoms? a. The oxygen atom in methanol b. The nitrogen atom in trimethylamine c. The phosphorus atom in :PH3 C. Molecular Orbital Theory Molecular Orbital Theory is one of two models that describes covalent bond formation. describes a covalent bond as resulting from combination of atomic orbitals to give molecular orbitals, which belong to the entire molecule • Molecular Orbital (MO) - is a region where electrons are most likely to be found in a molecule. • It has a specific energy and general shape • There are two types of orbital combination: • Additive combination • Substractive combination • Additive combination: (bonding) MO is lower in energy • Substractive combination: (antibonding) MO is higher in energy Molecular Orbitals in Ethylene • The bonding MO is from combining p orbital lobes with the same algebraic sign • The antibonding MO is from combining lobes with opposite signs • Only bonding MO is occupied Chapter 1 Summary • Organic chemistry – chemistry of carbon compounds • Atom: positively charged nucleus surrounded by negatively charged electrons • Electronic structure of an atom described by wave equation – Electrons occupy orbitals around the nucleus. – Different orbitals have different energy levels and different shapes • s orbitals are spherical, p orbitals are dumbbellshaped Summary • 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 Summary • Sigma (s) bonds - Circular cross-section and are formed by head-on interaction • Pi () bonds – “dumbbell” shape from sideways interaction of p orbitals Summary • 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 Summary • Atoms such as nitrogen and oxygen hybridize to form strong, oriented bonds – The nitrogen atom in ammonia and the oxygen atom in water are sp3-hybridized