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

vancomycin C66H75Cl2N9O24

vancomycin C66H75Cl2N9O24

Compounds you may know

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

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!!

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

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®) 2nd point…ask what they put in cars and why it works Why does the alcohol burn? 9

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

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.

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-

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+

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

Group 13(B and Al) – forms 3+ Group 15 – forms 3- Group 16 – forms 2- 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!

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.

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)

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).

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.

1s, 2s, 3s

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

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

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

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

degenerate orbitals (same E)

Section 1.3 Electronic Configurations Good extra credit question: Why does Hund’s rule exist? B/c electrons repel each other, and placing them in the same orbital requires more energy than keeping them farther apart with parallel spins 28

Section 1.3 Electronic Configurations

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

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

4th is the highest Energy level, so there are 7 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

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

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).

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.

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).

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

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

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?

Electronegativity

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

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?

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

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)

HONC - 1234

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

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

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

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

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?

Let’s consider methane (CH4) How many valence electrons? 4 In which orbitals? 2s22p2 How many unpaired electrons (from orbital diagram)? 2 – indicating the possibility for 2 bonds But I thought carbon always formed 4 bonds??? So, both the 2s and 2p orbitals are used to form bonds How many bonds does carbon form? How do we explain this? Hybridization

Hybridization The s and p orbitals of the C atom combine with each other to form hybrid orbitals before they combine with orbitals of another atom to form a covalent bond. Lone pairs MUST be accounted for in hybridization of orbitals since they are also negatively charged regions causing repulsion (regions of electron density).

Formation of an sp3 hybrid C Mathematically rearranged… (promotion) Why does it get that shape? …then hybridization 1s 2sp3

Formation of an sp3 hybrid The sp3 hybrids now arrange themselves as far away as possible Tetrahedral shape Formation of CH4 comes from the 1s orbital of 4 hydrogens overlapping with the 2sp3 hybrids in carbon 1s 2sp3 These electrons from the 1s of hydrogen A head on overlap of the s and sp3 orbitals is called a s bond. How many s bonds are in methane?

ethane sp3 hybrid orbital diagram of ethane (C2H6) -draw the orbital diagrams for both C How many s bonds does ethane have? 1s 2sp3 1s 2sp3

Formation of an sp2 hybrid Mathematically rearranged… (promotion) …then hybridization Unstable, but only exists for a short period of time 1s 2p 2sp2

Formation of an sp2 hybrid In the formation of ethylene (C2H4), each carbon becomes sp2 hybridized 4 hydrogens come in and pair in the lowest energy orbitals available (the 2sp2) These electrons pair to make a p bond 2p 2p These electrons pair to make a s bond 2sp2 2sp2 1s 1s

Formation of an sp2 hybrid = A σ bond results from head-to-head overlap in orbitals (s-s, p-p, s-p, or hybrids) [σ = sigma] A π bond results from side-to-side overlap in orbitals (p-p) Net result in ethylene (ethene) is a double bond (one σ, one π) Double bonds are shorter than single bonds and stronger than single bonds.

Bond Length So, why are double bonds shorter than single bonds? Which orbital is closer to the nucleus, s or p? What is the %s character in an sp3 hybridized orbital? What is the %s character in an sp2 hybridized orbital? The one with more %s character will be closer to the nucleus, thus the bond between two of these orbitals will be closer.

Bond Strength So, why are double bonds higher in energy than single bonds? Which bond involves the higher energy orbitals p or s? How many electrons are involved in a double bond compared to a single bond? You would think that a double bond is twice the bond strength as a single bond, but due to the poor overlap of the p-orbitals (p bond), that overlap is not a strong as the s bonds. So, the double bond is less than double the strength of a single bond. 1s 2p 2sp2

Formation of an sp hybrid Mathematically rearranged… 1s 2p 2sp

Formation of an sp hybrid In the formation of acetylene (C2H2), each carbon becomes sp hybridized 2 hydrogens come in and pair in the lowest energy orbitals available (the 2sp) These electrons pair to make 2 p bonds 1s 2p 2sp 2p These electrons pair to make a s bond 2sp 1s

Formation of an sp hybrid A σ bond results from the s-p overlap from the hydrogens to the carbons, and the sp-sp overlap between carbons 2 π bonds results from side-to-side overlap in orbitals (p-p) Net result in acetylene (ethyne) is a triple bond (one σ, two π) Are triple bonds longer or shorter than double bonds? Higher/lower energy? Triple bonds are shorter than double bonds and higher energy

Summary of Hybridization Hybridization of C sp3 sp2 sp Example methane ethylene ethene acetylene ethyne # Groups bonded to C 4 3 2 Geometry Tetrahedral Trigonal planar Linear Bond angles (o) 109.5 ~120 ~180 Types of bonds to C 4s 3s, 1p 2s, 2p C-C bond length (pm) 154 134 120 C-C bond strength (kcal/mol) 90 174 231

Hybridizations, geometries and bond angles of compounds containing heteroatoms Ex. CH3NH2 Ex. CH3OH Ex. CH3COCH3

carbocation CH3+

carbanion CH3-

Condensed Structures

Structural Formulas Complete Condensed Structural formulas actually show the arrangement of atoms Complete Lewis structures Condensed Lewis structures

Structural Formulas When double or triple bonds are included, they are usually shown in the condensed formula.

Chemical Formulas - Summary Molecular Formula - C2H4O2 Empirical Formula – CH2O Expanded Lewis structures (Structural Formula or Kekulé Structure) Condensed Lewis structures - CH3COOH Line-angle structures (Skeletal structures)

Line-Angle Formulas Carbons aren’t always shown. Carbons are assumed to be present at angles/intersections. Hydrogen atoms bonded to carbon are not shown. There are assumed to be enough hydrogens to give carbon 4 bonds. Exception: aldehydes -CHO Heteroatoms (anything other than C and corresponding hydrogens) ARE shown.

Example: Proline

Example: Vitamin A

Line-Angle Formulas

Review question: Cocaine is shown below Review question: Cocaine is shown below. Determine the geometry, bond angles, and hybridizations at all highlighted locations (careful with the N). Also, provide a molecular formula for the compound.