1-1 Organic Chemistry William H. Brown Christopher S. Foote Brent L. Iverson William H. Brown Christopher S. Foote Brent L. Iverson

1-2 Covalent Bonding & Shapes of Molecules Chapter 1

1-3 Organic Chemistry  The study of the compounds of carbon  Over 10 million compounds have been identified about 1000 new ones are identified each day!  C is a small atom it forms single, double, and triple bonds it is intermediate in electronegativity (2.5) it forms strong bonds with C, H, O, N, and some metals

1-4 Schematic View of an Atom a small dense nucleus, diameter m, which contains positively charged protons and most of the mass of the atom an extranuclear space, diameter m, which contains negatively charged electrons

1-5 Electron Configuration of Atoms  Electrons are confined to regions of space called principle energy levels (shells) each shell can hold 2n 2 electrons (n = 1,2,3, )

1-6 Electron Configuration of Atoms  Shells are divided into subshells called orbitals, which are designated by the letters s, p, d, f, s (one per shell) p (set of three per shell 2 and higher) d (set of five per shell 3 and higher).....

1-7 Electron Configuration of Atoms  Aufbau Principle: orbitals fill in order of increasing energy from lowest energy to highest energy  Pauli Exclusion Principle: only two electrons can occupy an orbital and their spins must be paired  Hund’s Rule: when orbitals of equal energy are available but there are not enough electrons to fill all of them, one electron is added to each orbital before a second electron is added to any one of them

1-8 Electron Configuration of Atoms  The pairing of electron spins

1-9 Electron Configuration of Atoms  Table 1.3 The Ground-State Electron Configuration of Elements 1-18

1-10 Lewis Dot Structures  Gilbert N. Lewis  Valence shell: the outermost occupied electron shell of an atom  Valence electrons: electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions  Lewis dot structure: the symbol of an element represents the nucleus and all inner shell electrons dots represent valence electrons

1-11 Lewis Dot Structures  Table 1.4 Lewis Dot Structures for Elements 1-18

1-12 Lewis Model of Bonding  Atoms bond together so that each atom acquires an electron configuration the same as that of the noble gas nearest it in atomic number anionan atom that gains electrons becomes an anion cationan atom that loses electrons becomes a cation ionic solidsthe attraction of anions and cations leads to the formation of ionic solids covalent bondan atom may share electrons with one or more atoms to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bond polar covalent bondsbonds may be partially ionic or partially covalent; these bonds are called polar covalent bonds

1-13 Electronegativity  Electronegativity: a measure of an atom’s attraction for the electrons it shares with another atom in a chemical bond  Pauling scale generally increases left to right in a row generally increases bottom to top in a column

1-14 Formation of Ions  A rough guideline: ions will form if the difference in electronegativity between interacting atoms is 1.9 or greater example: sodium (EN 0.9) and fluorine (EN 4.0) we use a single-headed (barbed) curved arrow to show the transfer of one electron from Na to F in forming Na + F -, the single 3s electron from Na is transferred to the partially filled valence shell of F

1-15 Covalent Bonds  The simplest covalent bond is that in H 2 the single electrons from each atom combine to form an electron pair the shared pair functions in two ways simultaneously; it is shared by the two atoms and fills the valence shell of each atom  The number of shared pairs one shared pair forms a single bond two shared pairs form a double bond three shared pairs form a triple bond

1-16 Polar and Nonpolar Covalent Bonds  Although all covalent bonds involve sharing of electrons, they differ widely in the degree of sharing  We divide covalent bonds into nonpolar covalent bonds polar covalent bonds

1-17 Polar and Nonpolar Covalent Bonds an example of a polar covalent bond is that of H-Cl the difference in electronegativity between Cl and H is = 0.9  +  -we show polarity by using the symbols  + and  -, or by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end

1-18 Polar Covalent Bonds  Bond dipole moment (  ): a measure of the polarity of a covalent bond the product of the charge on either atom of a polar bond times the distance between the nuclei Table 1.7 shows average bond dipole moments of selected covalent bonds

1-19 Lewis Structures  To write a Lewis structure determine the number of valence electrons determine the arrangement of atoms connect the atoms by single bonds arrange the remaining electrons so that each atom has a complete valence shell show a bonding pair of electrons as a single line show a nonbonding pair of electrons as a pair of dots in a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons

1-20 Lewis Structures - Table 1.3  In neutral molecules hydrogen has one bond carbon has 4 bonds and no lone pairs nitrogen has 3 bonds and 1 lone pair oxygen has 2 bonds and 2 lone pairs halogens have 1 bond and 3 lone pairs

1-21 Formal Charge  Formal charge:  Formal charge: the charge on an atom in a molecule or a polyatomic ion  To derive formal charge 1. write a correct Lewis structure for the molecule or ion 2. assign each atom all its unshared (nonbonding) electrons and one-half its shared (bonding) electrons 3. compare this number with the number of valence electrons in the neutral, unbonded atom

1-22 Formal Charge  Example:  Example: Draw Lewis structures, and show which atom in each bears the formal charge

1-23 Exceptions to the Octet Rule  Molecules containing atoms of Group 3A elements, particularly boron and aluminum Aluminum chloride : : : FB F F ClAl Cl Cl 6 electrons in the valence shells of boron and aluminum Boron trifluoride :: :: :: : : : : : : : : :

1-24 Exceptions to the Octet Rule  Atoms of third-period elements have 3d orbitals and may expand their valence shells to contain more than 8 electrons phosphorus may have up to 10

1-25 Exceptions to the Octet Rule sulfur, another third-period element, forms compounds in which its valence shell contains 8, 10, or 12 electrons

1-26 Functional Groups  Functional group:  Functional group: an atom or group of atoms within a molecule that shows a characteristic set of physical and chemical properties  Functional groups are important for three reason; they are 1. the units by which we divide organic compounds into classes 2. the sites of characteristic chemical reactions 3. the basis for naming organic compounds

1-27 Alcohols hydroxyl  contain an -OH (hydroxyl) group condensed structural formula  Ethanol may also be written as a condensed structural formula

1-28 Alcohols alcohols are classified as primary (1°), secondary (2°), or tertiary (3°) depending on the number of carbon atoms bonded to the carbon bearing the -OH group

1-29 Alcohols there are two alcohols with molecular formula C 3 H 8 O

1-30 Amines amino group  contain an amino group; an sp 3 -hybridized nitrogen bonded to one, two, or three carbon atoms an amine may by 1°, 2°, or 3° CH 3 NH H CH 3 NH CH 3 CH 3 NCH 3 CH 3 Methylamine (a 1° amine) Dimethylamine (a 2° amine) Trimethylamine (a 3° amine) : : :

1-31 Aldehydes and Ketones carbonyl (C=O) group  contain a carbonyl (C=O) group

1-32 Carboxylic Acids carboxyl (-COOH) group  contain a carboxyl (-COOH) group

1-33 Carboxylic Esters  Ester:  Ester: a derivative of a carboxylic acid in which the carboxyl hydrogen is replaced by a carbon group

1-34 Carboxylic Amide arboxylic amide amide  Carboxylic amide, commonly referred to as an amide: a derivative of a carboxylic acid in which the -OH of the -COOH group is replaced by an amine the six atoms of the amide functional group lie in a plane with bond angles of approximately 120°

1-35 VSEPR  Based on the twin concepts that atoms are surrounded by regions of electron density regions of electron density repel each other

1-36 VSEPR Model  Example:  Example: predict all bond angles for these molecules and ions

1-37 Polar and Nonpolar Molecules  To determine if a molecule is polar, we need to determine if the molecule has polar bonds the arrangement of these bonds in space  Molecular dipole moment (  ):  Molecular dipole moment (  ): the vector sum of the individual bond dipole moments in a molecule reported in debyes (D)

1-38 Polar and Nonpolar Molecules  these molecules have polar bonds, but each has a zero dipole moment

1-39 Polar and Nonpolar Molecules  these molecules have polar bonds and are polar molecules

1-40 Polar and Nonpolar Molecules formaldehyde has polar bonds and is a polar molecule

1-41 Resonance  For many molecules and ions, no single Lewis structure provides a truly accurate representation

1-42 Resonance  Linus Pauling s many molecules and ions are best described by writing two or more Lewis structures individual Lewis structures are called contributing structures connect individual contributing structures by double- headed (resonance) arrows the molecule or ion is a hybrid of the various contributing structures

1-43 Resonance  Examples: equivalent contributing structures

1-44 Resonance  Curved arrow:  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 is a survival skill in organic chemistry learn it well!

1-45 Resonance  All contributing structures must 1. have the same number of valence electrons 2. obey the rules of covalent bonding no more than 2 electrons in the valence shell of H no more than 8 electrons in the valence shell of a 2nd period element 3rd period elements, such as P and S, may have up to 12 electrons in their valence shells 3. differ only in distribution of valence electrons; the position of all nuclei must be the same 4. have the same number of paired and unpaired electrons

1-46 Resonance  The carbonate ion, for example a hybrid of three equivalent contributing structures the negative charge is distributed equally among the three oxygens

1-47 Resonance  Preference 1:  Preference 1: filled valence shells structures in which all atoms have filled valence shells contribute more than those with one or more unfilled valence shells Greater contribution; both carbon and oxygen have complete valence shells Lesser contribution; carbon has only 6 electrons in its valence shell + + CCH 3 OCH 3 O H HC H H

1-48 Resonance  Preference 2:  Preference 2: maximum number of covalent bonds structures with a greater number of covalent bonds contribute more than those with fewer covalent bonds

1-49 Resonance  Preference 3:  Preference 3: least separation of unlike charge structures with separation of unlike charges contribute less than those with no charge separation Lesser contribution (separation of unlike charges) CH 3 -C-CH 3 CH 3 -C-CH 3 Greater contribution (no separation of unlike charges) O - O :: :::

1-50 Resonance  Preference 4:  Preference 4: negative charge on the more electronegative atom structures that carry a negative charge on the more electronegative atom contribute more than those with the negative charge on the less electronegative atom

1-51 Quantum or Wave Mechanics  Albert Einstein: E = h (energy is quantized ) light has particle properties  Louis deBroglie: wave/particle duality  Erwin Schrödinger: wave equation wave function, wave function,  : a solution to a set of equations that depicts the energy of an electron in an atom each wave function is associated with a unique set of quantum numbers orbitaleach wave function occupies three-dimensional space and is called an orbital  2  2 is the probability of finding an electron at a given point in space

1-52 Shapes of 1s and 2s Orbitals  Probability distribution (  2 ) for 1s and 2s orbitals showing an arbitrary boundary surface containing about 95% of the electron density

1-53 Shapes of a Set of 2p Atomic Orbitals  Three-dimensional shapes of 2p atomic orbitals

1-54 Molecular Orbital Theory  Electrons in atoms exist in atomic orbitals  Electrons in molecules exist in molecular orbitals (MOs)  Using the Schrödinger equation, we can calculate the shapes and energies of MOs

1-55 Molecular Orbital Theory  Rules: combination of n atomic orbitals (mathematically adding and subtracting wave functions) gives n MOs (new wave functions) MOs are arranged in order of increasing energy MO filling is governed by the same rules as for atomic orbitals: Aufbau principle: fill beginning with LUMO Pauli exclusion principle: no more than 2e - in a MO Hund’s rule: when two or more MOs of equivalent energy are available, add 1e - to each before filling any one of them with 2e -

1-56 Molecular Orbital Theory  Terminology ground state = lowest energy state excited state = NOT lowest energy state  = sigma bonding MO  * = sigma antibonding MO  = pi bonding MO  * = pi antibonding MO HOMO = highest occupied MO LUMO = lowest unoccupied MO

1-57 Molecular Orbital Theory  Sigma 1s bonding and antibonding MOs

1-58 Molecular Orbital Theory MO energy diagram for H 2 : (a) ground state and (b) lowest excited state

1-59 Molecular Orbitals computed sigma bonding and antibonding MOs for H 2

1-60 Molecular Orbitals  pi bonding and antibonding MOs

1-61 Molecular Orbitals computed pi bonding and antibonding MOs for ethylene

1-62 Molecular Orbitals computed pi bonding and antibonding orbitals for formaldehyde

1-63 Hybrid Orbitals  The Problem: bonding by 2s and 2p atomic orbitals would give bond angles of approximately 90° instead we observe bond angles of approximately 109.5°, 120°, and 180°  A Solution hybridization of atomic orbitals 2nd row elements use sp 3, sp 2, and sp hybrid orbitals for bonding

1-64 Hybrid Orbitals  Hybridization of orbitals (L. Pauling) the combination of two or more atomic orbitals forms a new set of atomic orbitals, called hybrid orbitals  We deal with three types of hybrid orbitals sp 3 sp 3 (one s orbital + three p orbitals) sp 2 sp 2 (one s orbital + two p orbitals) sp sp (one s orbital + one p orbital)  Overlap of hybrid orbitals can form two types of bonds depending on the geometry of overlap  bonds  bonds are formed by “direct” overlap  bonds  bonds are formed by “parallel” overlap

1-65 sp 3 Hybrid Orbitals each sp 3 hybrid orbital has two lobes of unequal size the sign of the wave function is positive in one lobe, negative in the other, and zero at the nucleus the four sp 3 hybrid orbitals are directed toward the corners of a regular tetrahedron at angles of 109.5°

1-66 sp 3 Hybrid Orbitals orbital overlap pictures of methane, ammonia, and water

1-67 sp 2 Hybrid Orbitals the axes of the three sp 2 hybrid orbitals lie in a plane and are directed toward the corners of an equilateral triangle the unhybridized 2p orbital lies perpendicular to the plane of the three hybrid orbitals

1-68 Bonding in Ethylene

1-69 Bonding in Formaldehyde

1-70 sp Hybrid Orbitals two lobes of unequal size at an angle of 180° the unhybridized 2p orbitals are perpendicular to each other and to the line created by the axes of the two sp hybrid orbitals

1-71 Bonding in Acetylene, C 2 H 2

1-72 Hybrid Orbitals

1-73 Bond Lengths and Bond Strengths

1-74 Covalent Bonds & Shapes of Molecules End Chapter 1