The place to begin a study of Organic chemistry is with chemical bonds of carbon-to-carbon, carbon to hydrogen, or carbon to other atoms.

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1 1 Chapter 3. Bonding The most fundamental concept in Organic chemistry is the nature of the bond between two carbon atoms or between carbon and another atom. For the most part, these are covalent bonds to carbon, although ionic bonds will be seen. Most common organic molecules are characterized by the presence of covalent bonds. The place to begin a study of Organic chemistry is with chemical bonds of carbon-to-carbon, carbon to hydrogen, or carbon to other atoms.

2 To begin, you should know: 2 The electronic configuration of elements in the first two rows of the periodic table. The shape of s- and p- atomic orbitals and how they relate to electronic configuration. The difference among s- and p- and d- orbitals. The difference between an ionic and a covalent bond. A sense of difference in the size of the elements and their respective ions. Covalent bonds are made of shared electrons. The concept of electronegativity.

3 When completed, you should know: 3 s-orbitals are spherically symmetrical, p-orbitals are "dumbbell" shaped, and hybrid orbitals are directional. Electrons are found in atomic orbitals at discrete distances from the nucleus in an atom. Electronic configuration of an atom describes the configuration associated with electrons in atomic orbitals. Electrons in the bond of a molecule are located between two nuclei and are at different energy levels than in an unbonded atom. Ionic bonds are formed by the electrostatic attraction of two atoms or groups that have opposite charges. Covalent bonds are made of two electrons that are mutually shared between two atoms. Mixing atomic orbitals forms hybrid molecular orbitals (LCAO method); s- and p-orbitals can be mixed to form the hybrid, which determines the hybridization (e.g., sp 3 ). Organic molecules generally have a backbone of carbon carbon covalent bonds. Polarized bonds are formed when two atoms are bonded together, but one is more electronegative. Polarized covalent bonds are generally weaker than nonpolarized covalent bonds. The VSEPR model is used to predict the three-dimensional shape around an individual atom. Reactions are driven by making and breaking bonds, which releases or requires energy.

4 Electrons are Important 4 Atoms are discreet entities that differ from one another by the number of protons, neutrons and electrons that make up each atom. Protons and neutrons are found in the nucleus, of course, but reactions involving organic molecules do not involve transfer, gain or loss of protons or neutrons. Chemical reactions involve the transfer of electrons, which are the important non-nuclear constituents of an atom. To determine chemical reactivity, a method has been developed to ascertain the position of electrons relative to the nucleus. Electrons are said to reside in orbitals.

5 Electrons, Waves and Orbitals 5 The motion of electrons has some characteristics of wave motion. The motion of an electron is expressed by a wave equation, which has a series of solutions and each solution is called a wavefunction. Each electron may be a described by a wavefunction whose magnitude varies from point to point in space. A particular solution to the Schrödinger wave equation, for a given type of electron, is determined by the equation. Hψ = Eψ H is a mathematical operator called the Hamiltonian operator E is the numerical value for the energy ψ is a particular wavefunction. The relationship between orbitals and the Schrödinger equation is apparent when its solutions are represented as the waves shown in the Figure for various values of ψ that correspond to different energies. The amplitude of the wave is the wave function (ψ) and it has a Maximum (represented by +) and a minimum (represented by -), and each point in space can be represented by spatial coordinates (x,y,z). The point at which the wave change its phase is referred to as a node.

6 Electrons, Waves and Orbitals 6 Maximum amplitude is represented by + and a minimum amplitude is represented by - node d orbital (2 nodes) p orbital (1 node) (a) s orbital (0 nodes) Plotting ψ versus distance from the nucleus in the figure leads to the familiar s, p and d orbitals. (b) node

7 node Orbitals 7 d orbital (2 nodes) p orbital (1 node) s orbital (0 nodes) (a) (b) Orbitals can be viewed as charge clouds that represent points in space where electrons may be found node Plotting ψ versus distance from the nucleus in the figure leads to the familiar s, p and d orbitals. This small volume can be viewed as a charge cloud if it contains an electron, and the charge cloud represents the region of space where we are most likely to find the electron in terms of the (x,y,z) coordinates. The Heisenberg uncertainty principle states that the position and momentum of an electron cannot be simultaneously specified. It is only possible to determine the probability that an electron will be found at a particular point relative to the nucleus. Since the exact position of the electron is unknown (there is uncertainty as to its position), the probability of finding the electron in a unit volume of threedimensional space is given by /ψ(x,y,z)/ 2. The position is expressed as ψ/ 2 dt, which is the probability of an electron being in a small element of the volume δt.

8 Atomic Orbitals 8 s, p, and d orbitals s p d (4 types) d z 2 The energy levels represent the region in space where electrons are found relative to the nucleus. There are several quantum levels in atoms, particularly in high atomic mass elements. This means that there are different energy levels associated with each type of electron shell, so there are different types of s orbitals: 1s, 2s, 3s, 4s, etc., similar in shape but differing in energy (their relative distance from the nucleus). There are 2p, 3p, 4p orbitals, and 3d, 4d, 5d, and, 4f, 5f, 6f orbitals. Based on the Periodic Table: The first row elements (H, He) have only the spherical s-orbitals. The second row (Li, Be, B, C, N, O, F, He) has the 1s orbital and the 2s- and 2porbitals are in the outermost shell. The third row introduces 3s-, and 3p-orbitals. d-orbitals appear in the fourth row. Each shell will have one s, three p, five d and seven f orbitals (1, 2, 3, 4). * Note that d- and f-orbitals accept more electrons or give up more electrons in chemical reactions than a p-orbital because more orbitals are involved.

9 Electronic Configuration - 2nd Row 9 Electron distribution in atomic orbitals of an element is known as its electronic configuration. Each individual orbital can hold no more than two electrons. Electrons have the property of spin, which is associated with a magnetic dipole. The spin quantum number: selfrotation of the electron gives rise to an angular momentum vector. Each electron will have spin and the symbol is used to indicate an electron with a certain spin quantum number. A single orbital containing two electrons (a filled orbital) is represented by two opposed arrows ( ). 1 Li 2 Be 13 B 14 C 15 N 16 O 17 F 18 Ne H He 1s 2s 2p The symbol indicates that when an orbital contains two electrons, those two electrons are spin paired. Note that if two electrons occupy one orbital, spin pairing is lower in energy than if two electrons of the same spin are forced to occupy the same orbital. The Pauli exclusion principle states that if there are several orbitals of equal energy (such as the three 2p-orbitals), each orbital will fill with one electron before any orbitals contain two. 1s 1s

10 Electronic Configuration and the Aufbau Principle 10 Electrons have like charges. 2 electrons will repel in the same orbital. Adding two electrons to two 2p orbitals is lower in energy than adding two electrons to a single 2p orbital, so electrons "fill" orbitals to conserve energy. The concept of filling orbitals with electrons in ascending order of orbital energy until all available electrons have been used is known as the aufbau procedure. The order in which electrons fill is generalized by the mnemonic 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p Hydrogen has an electronic configuration 1s 1 and helium is 1s 2. The "1" represents the row of the Periodic Table, the "s" represents the orbital, and the superscript "1" represents the number of electrons in that orbital. Lithium begins the second row and the 2s orbitals begin to fill. The electron configuration of lithium is 1s 2 2s 1. This continues to the Noble gas neon with a configuration 1s 2 2s 2 2p 6. In the third row, sodium begins to fill the 3s orbital (1s 2 2s 2 2p 6 3s 1 ) and continues to argon, with a configuration of 1s 2 2s 2 2p 6 3s 2 3p 6.

11 1s 2p z 2p x 2p y 2s The nucleus is the convergence point of the tri-coordinate system. The orbitals are closely associated with the atomic nucleus (atomic orbitals), so the electrons are closely associated with the atom. An Atom in the Second Row 11 To construct an atom, there is a 1s orbital, a 2s orbital, and three 2p orbitals in the second row of equal energy The three p-orbitals have the same shape and energy. The space volume for each electron is described in a Cartesian (three-coordinate) system, x,y,z. The 1s orbital is represented by the spherical "green dot" at the center. The 2s orbital is represented by the "black circle" (meant to represent a sphere) that is larger than the 1s sphere, showing its greater distance from the nucleus. The three 2p orbitals are labeled as red for the 2p z, blue for the 2p x, and yellow for the 2p y orbitals. All have same energy = degenerate The purpose of this simplistic picture is to give a mental image that a 2p electron is more easily removed than the 1s or 2s electrons because it is further from the nucleus.

12 Chemical Bonds 12 A chemical reaction between atoms or groups of atoms will usually produce new combinations of atoms in a MOLECULE, held together by what is called a chemical bond. When one element reacts with another it does so via its electrons, not by the protons and neutrons in the nucleus, and the resulting bond between the atoms is composed of two electrons. Two major types of bonds will be considered. A covalent bond is formed by the mutual sharing of TWO valence electrons between two atoms. In other words, sharing electron density holds the atoms together. An ionic bond is formed transfer of electrons from one atom to another, resulting in ions (+ and ) that are held together by electrostatic attraction.

13 Ionic and Covalent Bonds 13 An ionic bond is formed transfer of electrons from one atom to another, resulting in ions (+ and ) that are held together by electrostatic attraction. As a practical matter, a molecule composed of two atoms on opposite sides of the periodic table Li and F, Na and Br, etc. tend to be ionic. A molecule formed by breaking a bond to generate a + or a - charge will be ionic since it will have a counterion with the opposite charge. A bond formed between two atoms in the middle of the periodic table will tend to be covalent

14 Ionization Potential and Electron Affinity 14 An atom with a valence electronic configuration such as Li (Group 1) will lose an electron during an electron transfer process, and an atom with a configuration such as fluorine (Group 17) will gain one electron. The energy required for the loss of one electron from an atom is called its ionization potential. The energy required for the gain of one electron into an atom is called its electron affinity. This means that the energy gained or lost for an atom is a measurable quantity. Electron transfer to form ions is the basis for the known ionic bonding in many molecules composed of alkali metals (Groups 1 and 2) and halogens (in Group 17): LiF, NaCl, KBr, NaI, etc.

15 Forming Ions 15 LiF has an ionic bond, where the positively charged Li is electrostatically bound to the negatively charged F. The "octet rule" is a useful tool, and it states that in the second row, a maximum of eight electrons can occupy the valence shell (the Ne configuration). In the first row, a maximum of two electrons can occupy the valence shell (the He configuration) and the second row can accommodate a maximum of eight electrons to give the Ne configuration (a total of 10 electrons). Assume there is an energetic preference for transferring electrons to attain the Noble gas configurations. Loss of an electron from Li does not lead to helium, but to a positively charged lithium ion (Li + ) with a 1s 2 configuration. If electron transfer removes an electron from F to give F +, the electronic configuration is 1s 2 2s 2 2p 4. If electron transfer adds one electron to a fluorine atom, however, the result is F which has the 1s 2 2s 2 2p 6 configuration (the Ne configuration). Addition of an electron to fluorine does not give Ne, but rather the fluoride ion F with the configuration 1s 2 2s 2 2p 6. F is more stable than F +.

16 Ionic versus Covalent. Dissociation energy Temp to dissociate: >3500 K 3227 C Surface of the sun: 5505 C Temp to dissociate: >600 K 327 C (begins dissociation) H Li F H C H Ionic Group 1,2 with Group 17 Electrons donated to one Atom or the other - charged Ionization potential - E to donate electron Electron affinity - E to accept electron H C in middle Shares electrons Covalent bond neutral

17 Bond Disruption and Bond Strength 17 For NaCl, the bond dissociation energy of the Na Cl bond is formally reported to be 98 kcal mol -1 and represents the amount of energy released when that bond is broken. In an organic compound, the bond dissociation energy of the C C bond is reported to be 145 kcal mol -1. The numbers suggest that the C-C bond is more difficult to break than the NaCl. This is not true! Most C-C bonds are completely disrupted at temperatures between C. The covalent bond is potentially much weaker than an ionic bond in terms of bond dissociation energy - not necessarily in terms of a reaction. In a reaction, disruption of the bond is not just "ripping" the two atoms apart, but rather a chemical process where electrons are transferred leading to the bond being broken

18 Breaking Bonds in a Reaction 18 In a reaction, the bond is not "ripping" the two atoms apart, but rather a chemical process where electrons are transferred from one atom to another to form a bond, leading to another bond being broken A Chemical Reaction Bond broken Bond formed A-B + C A + B-C Reactants Products If the A-B bond is such that A can accept electron density (from C), AND if C is an electron donor - then C can attack (donate electrons to) B, forming a new B-C bond, and breaking the A-B bond. In this reaction, the A-B bond is weak in that the reaction can occur - not the same thing as bond dissociation energy. Remember the acid-base reaction H H O H Cl H H O H Cl

19 Bond Strength Measured by a Chemical Reaction Bond dissociation energy is the energy required to completely disrupt the bond holding two atoms together. Bond dissociation energy is not necessarily the best measure of bond strength for a reaction. For the most part, bond strength in this course will refer to the ease of breaking the bond in a chemical reaction. There is a difference in separating ions in an ionic bond and breaking a bond in a covalent molecule that will be discussed later. For the moment, simply be able to recognize an ionic bond and a covalent bond.

20 Carbon 20 Carbon has an electronic configuration of 1s 2 2s 2 2p 2 with four valence electrons. The bonds in molecules that contain carbon are usually formed by sharing electrons with another atom in what is known as a covalent bond, rather than by complete transfer of electrons to form an ionic bond. Carbon has the electronic configuration 1s 2 2s 2 2p 2, so there are four electrons in the outmost shell. To achieve the helium configuration by electron transfer (1s 2 ), four electrons must be lost (fourth ionization potential). To achieve the Neon configuration (1s 2 2s 2 2p 6 ) four electrons must be gained (fourth electron affinity). In both cases, the energy requirements are absurd. It is unlikely that an elemental carbon atom will form bonds by the same type of electron transfer as found with LiF. For reactions in this book, elemental carbon such as that found in diamond or graphite will not be converted into carbon containing molecules. To say that carbon "transfers electrons and shares electrons" does not mean that atomic carbon will react with other atoms or other compounds for form a bond to become covalent.

21 Electrons are found in orbitals, Carbon is found in Molecules 21 The orbitals for atomic carbon are different from the orbitals for covalent carbon in a molecule; carbon in the element is different from carbon in a molecule. The distinction is that the bonding in organic molecules is covalent and it involves carbon, but elemental carbon is not readily converted to such compounds. This means the direct conversion of elemental carbon (graphite or diamond for example) into an organic molecule such as methane is unlikely. Both graphite and diamond are known for their stability, not for their ability to react with things. The fundamental conclusion is that the energy and position of electrons in molecular carbon is different from that in elemental (atomic) carbon.

22 Covalent Carbon-Carbon Bond 22 When electrons are concentrated in orbitals on a single atom, the electrons are said to be "localized" on that atom. When an atom is part of a covalent bond, it shares electrons with the other atom; i.e., the electrons in the bond are not localized on one atom but rather are shared by both atoms and concentrated between the atoms. The electrons that constitute the bond are concentrated between both nuclei rather than localized on one or the other. nucleus-1 nucleus-2 Two atoms that are not bonded. The electrons in each p orbital are localized on the atom As the atoms are brought within bonding distance, the orbitals of one atom must be directed towards the orbital of the other in order to share the electrons to form a covlaent bond maximum orbital overlap (maximum electron density) nucleus nucleus When the covalent bond is formed, the orbital on each atom has been distorted - the orbital is directed towards the other atom. The electrons are now shared between the atoms and the greatest concentration of electron density is between the nuclei (black atoms)

23 Sigma Bond (σ-bond) 23 maximum orbital overlap (maximum electron density) nucleus nucleus The orbitals used to form covalent bonds in a molecule are called molecular orbitals (they are different from atomic orbitals). This means that the electrons in the covalent bond are found in molecular orbitals. When two identical atoms share electrons in a covalent bond, most (but not all) of the electron density is distributed between the two nuclei (in the "space" between the two atoms). The electron density that constitutes the bond is located in orbitals between the nuclei. This means that the charge density of the shared electrons is greatest between the two nuclei. This orbital picture describes a sigma bond (σ bond).