Relevance to Organic Chemistry H H2C

CH2 +

ethylene

H–Cl

ether

Cl–

H2C CH2 carbocation

1

H H2C O + formaldehyde

H–Cl

–

ether

H2C O Cl oxocarbenium ion

If an organic molecule can donate electrons, it may react as a base

• Modify thinking as to what constitutes a weak or strong acid or base. • In ethylene there are two bonds between adjacent carbon atoms, a double bond (C=C), a strong (the σ-bond) and a weak π-bond. • When ethylene is mixed with the strong acid HCl, the weaker π-bond reacts with the HCl by donating two electrons from the weak π-bond to form a new CH bond and generate a positive charge on the adjacent carbon (a carbocation), with a chloride counterion. • Formaldehyde also has a strong σ-bond and a weak π-bond, this time in a carbon-oxygen double bond (C=O). • Formaldehyde reacts with HCl to give an oxocarbenium ion, the conjugate acid and chloride ion is the conjugate base. • It is clear that ethylene and formaldehyde react as weak bases with the strong acid HCl to give the indicated products.

Relevance to Organic Chemistry

• Categorizing reactions by their ability to donate electrons from one source to another, or from one atom or group to another.

• A focus on the bond making-bond breaking process is a key to understanding how reactions work.

2

Define Nucleophiles

• If a species donates two electrons to an atom such as B or Al, it is called a Lewis base. • In some cases, a molecule, a compound or an ion can donate two electrons to a carbon atom, forming a new bond to that carbon. • A special name is given to a species that donates electrons to carbon. It is called a nucleophile. • A nucleophile is essentially a Lewis base-type species that reacts with a carbon species that functions similar to a Lewis acid, but it does NOT form an ate complex • With an acid-base theme, a nucleophile can be thouight of as a modified Lewis base and the carbon it react with is a modified Lewis acid.

3

4

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.

To begin, you should know:

• The electronic configuration of elements in the first two rows. • The difference between s and p and d orbitals. • The shape of s and p atomic orbitals and how they relate to electronic configuration. • The difference between and an ionic bond and a covalent bond. • A sense of difference in the size of the elements and their respective ions. • Bonds are made of electrons. • The concept of electronegativity.

5

When completed, you should know: • s-Orbitals are spherically symmetrical and p-orbitals are "dumbbell" shaped. • Electrons are found in orbitals at discrete distances from the nucleus in an atom. • 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 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, and the number of s and p orbitals used to form the hybrid determines the hybridization (sp3 for example). • Organic molecules generally have a backbone of carbon-carbon covalent bonds. • Electronegativity for an atom increase to the right and up the Periodic Table. • Polarized bonds are formed when two atoms are bonded together but one is more electronegative. • Polarized covalent bonds are generally weaker than non-polarized covalent bonds. • Reactions are driven by making and breaking bonds, which release or require energy. • The VSEPR model is used to predict the three-dimensional shape around an individual atom.

6

Electrons are Important

• 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.

7

Electrons, Waves and Orbitals • 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. Hy = Ey node • H is a mathematical operator called the Hamiltonian operator • E is the numerical value for the energy • y is a particular wavefunction. • The relationship between orbitals d orbital (2 nodes) and the Schrödinger equation is apparent when its solutions are (a) represented as the waves shown in the Figure for various p orbital (1 node) values of y that correspond to different energies. • The amplitude of the wave is the s orbital (0 nodes) wave function (y) and it has a maximum (represented by +) and a minimum (represented by –), and each point in space can be represented by (b) spatial coordinates (x,y,z). • The point at which the wave changes Its phase is referred to as a node. node

8

node

Orbitals

9

d orbital (2 nodes)

These charge clouds are known as orbitals.

(a) p orbital (1 node) s orbital (0 nodes)

(b)

node

• 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 three-dimensional space is given by /y(x,y,z)/2. The position is expressed as ⁄y/2dt, which is the probability of an electron being in a small element of the volume dt. • 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.

• Plotting y versus distance from the nucleus in the Figure leads to the familiar s, p and d orbitals.

10

Atomic Orbitals s, p, and d orbitals

s

p

d (4 types)

d z2

• 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.

node

Orbitals

11

d orbital (2 nodes)

These charge clouds are known as orbitals.

(a) p orbital (1 node) s orbital (0 nodes)

(b)

node

• 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 three-dimensional space is given by /y(x,y,z)/2. The position is expressed as ⁄y/2dt, which is the probability of an electron being in a small element of the volume dt. • 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.

• Plotting y versus distance from the nucleus in the Figure leads to the familiar s, p and d orbitals.

12

Atomic Orbitals s, p, and d orbitals

s

p

d (4 types)

d z2

• 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.

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Electronic Configuration - 2nd Row 1s

• 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.

1s

H

1s

He

2s

2p

1 Li 2 Be 13 B 14 C

• Each electron will have spin and the symbol ↑ is used to indicate an electron with a certain spin quantum number.

15 N

• A single orbital containing two electrons (a filled orbital) is represented by two opposed arrows (↑↓).

17 F

16 O

18 Ne

• 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.

Electronic Configuration and the Aufbau Principle • 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 1s1 and helium is 1s2. • 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 1s22s1. • This continues to the Noble gas neon with a configuration 1s22s22p6. • In the third row, sodium begins to fill the 3s orbital (1s22s22p63s1) and continues to argon, with a configuration of 1s22s22p63s23p6.

14

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

Chemical Bonds • 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.

16

Ionic and Covalent Bonds

•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

17

Ionization Potential and Electron Affinity

• 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.

18

Forming Ions

19

• 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 1s2 configuration. • If electron transfer removes an electron from F to give F+, the electronic configuration is 1s22s22p4. • If electron transfer adds one electron to a fluorine atom, however, the result is F– which has the 1s22s22p6 configuration (the Ne configuration). • Addition of an electron to fluorine does not give Ne, but rather the fluoride ion F– with the configuration 1s22s22p6. F– is more stable than F+.