Introduction to Organic chemistry NOTES: Chapter 1

Chapter 1: Covalent Bonding and Shapes of Molecules

Organic chemistry is the study of the compounds of carbon.
organic compounds are everywhere around us.
The organic chemistry is the chemistry of carbon and only a few other elements—chiefly hydrogen, oxygen, and nitrogen. Many of organic compounds also contain sulfur, phosphorus, and a halogens.

1.1 How Do We Describe the Electronic Structure of Atoms?
The atom contains a small, dense nucleus made of neutrons and positively charged protons.
Electrons move around a nucleus in confined regions of space called principal energy levels or, shells. We number these shells 1, 2, 3, and so forth from the inside out.
Shells are divided into subshells designated by the letters s, p, d, and f, and within these subshells, electrons are grouped in orbitals.
An orbital is a region of space that can hold 2 electrons.

A. Electron Configuration of Atoms
The electron configuration of an atom is a description of the orbitals the electrons in the atom occupy.
We determine the ground-state electron configuration of an atom with the use of the following three rules:
Rule 1. Orbitals fill in order of increasing energy from lowest to highest
Rule 2. Each orbital can hold up to two electrons, each electron spins in a direction opposite that of its partner.
Rule 3. When orbitals of equivalent energy are available, but there are not enough electrons to fill them completely, then we add one electron to each equivalent orbital before we add a second electron to any one of them.

B. Lewis Structures
Chemists often focus on the outermost shell of the atoms, because electrons in this shell are the ones involved in the formation of chemical bonds and in chemical reactions.
The outer-shell electrons known as valence electrons, and the energy level in which they are found is known as the valence shell.
The electron configuration of carbon is 1s22s22p2, and has four valence (outer-shell) electrons.
To represent the outermost electrons of an atom, we use a representation called a Lewis structure.
A Lewis structure shows the symbol of the element, surrounded by a number of dots equal to the number of electrons in the outer shell of an atom of that element.
In Lewis structures, the atomic symbol represents the nucleus and all filled inner shells.
Table 1.3 shows Lewis structures for the first 18 elements of the Periodic Table.
The number of valence electrons of the element corresponds to the group number of the element in the Periodic Table; for example, oxygen, with six valence electrons, is in Group 6A. (The the exception of that rule is helium).
Because of the differences in number and kind of valence shell orbitals available to elements of the second and third periods, significant differences exist in the covalent bonding of oxygen and sulfur and of nitrogen and phosphorus.

1.2 What Is the Lewis Model of Bonding?
A. Formation of Ions
Lewis pointed out that the chemical inertness of the noble gases indicates a high degree of stability of the electron configurations.
The atoms tends to react in ways that achieve an outer shell of eight valence electrons, and this tendency is known as the octet rule.
An atom with almost eight valence electrons tends to gain the needed electrons to have eight electrons in its valence shell and an electron configuration like that of the noble gas nearest it in atomic number.
When the atom gains electron it becomes a negatively charged ion and called an anion.
An atom with only one or two valence electrons tends to lose the number of electrons required to have the same electron configuration as the noble gas nearest it in atomic number.
In losing one or more electrons, the atom becomes a positively charged ion called a cation.

B. Formation of Chemical Bonds
The atoms interact with each other in such a way that each atom participating in a chemical bond acquires a valence-shell electron configuration the same as that of the noble gas closest to it in atomic number.
Atoms acquire completed valence shells in two ways:
1. An atom may lose or gain enough electrons to acquire a filled valence shell. An atom that gains electrons becomes an anion, and an atom that loses electrons becomes a cation. A chemical bond between an anion and a cation is called an ionic bond.
2. An atom may share electrons with one or more other atoms to acquire a filled valence shell. A chemical bond formed by sharing electrons is called a covalent bond.
Ionic bonds usually form between a metal and a nonmetal. For example reaction of metal sodium with a nonmetal chlorine to form sodium chloride, Na⁺Cl⁻.
When two nonmetals or a metalloid and a nonmetal combine, the bond between them is usually covalent. For example for two nonmetals Cl₂, H₂O, CH₄, and NH₃. Examples for a metalloid and a nonmetal include BF₃, SiCl₄, and AsH₄.

C. Electronegativity and Chemical Bonds
Electronegativity is a measure of the force of an atom’s attraction for electrons that it shares in a chemical bond with another atom.
The most widely used scale of electronegativities shown in table.
The fluorine is the most electronegative element.
The electronegativity values increases from left to right within a period of the Periodic Table because of the increasing positive charge on the nucleus, which leads to a stronger attraction for electrons in the valence shell, and generally decreasing from top to bottom within a group because of the increasing distance of the valence electrons from the nucleus, which leads to weak attraction between a nucleus and its valence electrons.
The electronegativity of a particular element depends also on its oxidation state. For example the electronegativity of Cu(I) in Cu₂O, for example, is 1.8, whereas of Cu(II) in CuO is 2.0.

Ionic Bonds
Forms by the transfer of electrons from the valence shell of an atom of lower electronegativity to the valence shell of an atom of higher electronegativity.
The more electronegative atom gains one or more valence electrons and becomes an anion; the less electronegative atom loses one or more valence electrons and becomes a cation.
This type of electron transfer is occur if the difference in electronegativity between two atoms is approximately 1.9 or greater.
For example the bond that formed between sodium (electronegativity 0.9) and fluorine (electronegativity 4.0) is ionic. The difference in electronegativity between these two elements is 3.1. In forming Na⁺F⁻, the single 3s valence electron of sodium is transferred to the partially filled valence shell of fluorine and hence both sodium and fluorine form ions that have the same electron configuration as neon, the noble gas closest to each in atomic number.

Covalent Bonds
A covalent bond forms when electron pairs are shared between two atoms whose difference in electronegativity is 1.9 or less. For example a covalent bond is that in a hydrogen molecule, H₂ .
When two hydrogen atoms bond, the single electron from each atom combine to form an electron pair with the release of energy. A bond formed by sharing a pair of electrons is called a single bond and is represented by a single line between the two atoms.
The electron pair shared between the two hydrogen atoms in H₂ completes the valence shell of each hydrogen.
Thus, in H₂ , each hydrogen has two electrons in its valence shell and an electron configuration like that of helium, the noble gas nearest to it in atomic number:
In forming a covalent bond, an electron pair occupies the region between two nuclei and serves to shield one positively charged nucleus from the repulsive force of the other positively charged nucleus. At the same time, an electron pair attracts both nuclei.
The distance between nuclei participating in a chemical bond is called a bond length.
Every covalent bond has a definite bond length.
The covalent bonds can be classified into two categories: nonpolar covalent and polar covalent.
In a nonpolar covalent bond, electrons are shared equally. In a polar covalent bond, they are shared unequally.
The guidelines in Table 1.5 (see original Book) will help to decide whether a given bond is more likely to be nonpolar covalent, polar covalent, or ionic.
ِAs a consequence of the unequal sharing of electrons in a polar covalent bond , the more electronegative atom gains a greater fraction of the shared electrons and acquires a partial negative charge (δ⁻), and the less electronegative atom has a lesser fraction of the shared electrons and acquires a partial positive charge (δ⁺), which leds to form a dipole (two poles).
We can display the polarity of a covalent bond by a type of molecular model called an electron density model.
By studying many organic compounds, we can make the following generalizations: In neutral (uncharged) organic compounds:
1. H has one bond.
2. C has four bonds.
3. N has three bonds and one unshared pair of electrons.
4. O has two bonds and two unshared pair of electrons.
5. F, Cl, Br, and I have one bond and three unshared pairs of electrons.

E. Formal Charge
The charge on an atom in a molecule or polyatomic ion is called its formal charge.
To derive a formal charge:
Step 1: Write a correct Lewis structure for the molecule or ion.
Step 2: Assign to each atom all its unshared (nonbonding) electrons and one-half its shared (bonding) electrons.
Step 3: Compare the number arrived at in Step 2 with the number of valence electrons in the neutral, unbonded atom.

The elements of the second period, including carbon, nitrogen, and oxygen, can accommodate no more than eight electrons in the four orbitals (2s, 2px , 2py , and 2pz) of their valence shells.

1.3 How Do We Predict Bond Angles and the Shapes of Molecules?
We can predict bond angles in the molecules by using the concept of valence-shell electron-pair repulsion (VSEPR).
The VSEPR can be used to predict the shapes of molecules.

1.4 How Do We Predict If a Molecule
Is Polar or Nonpolar?
A molecule will be polar if (1) it has polar bonds and (2) the vector sum of its bond dipoles is not zero (i.e., the bond dipoles cancel each other).

1.5 What Is Resonance?
For a great many molecules and ions, no single Lewis structure provides a truly accurate representation.

A. The Theory of Resonance
Many molecules and ions are best described by writing two or more Lewis structures and considering the real molecule or ion to be a composite of these structures.
Individual Lewis structures known as resonance contributing structures.
The real molecule or ion is a resonance hybrid of the various contributing structures by interconnecting them with double-headed arrows.

B. Curved Arrows and Electron Pushing
To show how redistribution of valence electrons occurs, chemists use a symbol called a curved arrow, which shows the repositioning of an electron pair from its origin (the tail of the arrow) to its destination (the head of the arrow).
A curved arrow is symbol for keeping track of electron pairs or, as some call it, electron pushing.

C. Rules for Writing Acceptable Resonance Contributing Structures
You must follow these four rules in writing acceptable resonance contributing structures:
1. All contributing structures must have the same number of valence electrons.
2. All contributing structures must obey the rules of covalent bonding; thus, no contributing structure may have more than 2 electrons in the valence shell of hydrogen or more than 8 electrons in the valence shell of a second-period element. Third-period elements, such as sulfur and phosphorus, may have up to 12 electrons in their valence shells.
3. The positions of all nuclei must be the same; that is, contributing structures differ only in the distribution of valence electrons.
4. All contributing structures must have the same total number of paired and unpaired electrons.

1.6 What Is the Orbital Overlap Model of Covalent Bonding?
Lewis and VSEPR models help us to understand covalent bonding and the geometry of molecules, but failed to explain the different in chemical reactivity for some bonds.
The formation of covalent bonds by the overlap of atomic orbitals helps us to understand the difference in chemical reactivity.

A. Shapes of Atomic Orbitals
By drawing a boundary surface around the region of space that encompasses some arbitrary percentage of the negative charge associated with that orbital, we'll find the s orbitalsfor example have the shape of a sphere with its center at the nucleus and the 1s orbital is the smallest, 2s orbital is a larger sphere, and a 3s orbital is an even larger sphere.
Each 2p orbital consists of two lobes arranged in a straight line with the nucleus in the middle.

B. Formation of a Covalent Bond by the Overlap of Atomic Orbitals
covalent bond is formed when a portion of an atomic orbital of one atom overlaps a portion of an atomic orbital of another atom.
For example in H2 two hydrogens approach each other so that their 1s atomic orbitals overlap to form a sigma (σ) covalent bond .
A sigma (σ) bond is a covalent bond in which orbitals overlap along the axis joining the two nuclei.

C. Hybridization of Atomic Orbitals
The atomic orbitals combine to form new orbitals, called hybrid orbitals.
The hybrid orbitals that result from the combination of s and p atomic orbitals named as the sp-type hybrid orbitals, which are three types sp3, sp2, and sp.

D. sp3 Hybrid Orbitals: Bond Angles of Approximately 109.5°
The combination of one 2s atomic orbital and three 2p atomic orbitals forms four equivalent sp3 hybrid orbitals.
The central atom in the compound uses four sp3 hybrid orbitals to either form a sigma (σ) bond with a hydrogen atom or to hold unshared pairs of electrons.

E. sp2 Hybrid Orbitals: Bond Angles of Approximately 120°
The combination of one 2s atomic orbital and two 2p atomic orbitals forms three equivalent sp2 hybrid orbitals.
The third 2p atomic orbital is not involved in hybridization and consists of two lobes lying perpendicular to the plane of the hybrid orbitals.
When the sp2 hybrid orbitals are overlapped along a common axis a sigma (σ) bond is formed. The remaining 2p orbitals which lie parallel to each other , they overlap to form a pi (π) bond.
A pi (π) bond is a covalent bond formed by the overlap of parallel p orbitals.
Pi bonds are generally weaker than sigma bonds.

F. sp Hybrid Orbitals: Bond Angles of Approximately 180°
The combination of one 2s atomic orbital and one 2p atomic orbital forms two equivalent sp hybrid orbitals.
Table 1.8 summarizes the relationship among the number of groups bonded to carbon, orbital hybridization, and the types of bonds involved.

1.7 What Are Functional Groups?
Functional groups are characteristic structural units by which we divide organic compounds into classes.
Functional groups:
- are sites of chemical reaction; a particular functional group, in whatever compound we find it, undergoes the same types of chemical reactions.
- determine, in large measure, the physical properties of a compound.

- are the units by which we divide organic compounds into families.
- serve as a basis for naming organic compounds.

A. Alcohols
The functional group of an alcohol is an ―OH (hydroxyl) group bonded to a tetrahedral carbon atom.
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.

B. Amines
The functional group of an amine is an amino group.
They are classified to a primary (1°) amine, nitrogen is bonded to one carbon atom, and secondary (2°) amine, it is bonded to two carbon atoms, and finally a tertiary (3°) amine, it is bonded to three carbon atoms.

C. Aldehydes and Ketones
Both aldehydes and ketones contain a C=O (carbonyl) group.
The aldehyde functional group contains a carbonyl group bonded to a hydrogen (―CHO). For example formaldehyde CH2O.
The functional group of a ketone is a carbonyl group bonded to two carbon atoms.

D. Carboxylic Acids
The functional group of a carboxylic acid is a _COOH (carboxyl: carbonyl + hydroxyl) group.