Chapter 1: Electronic Structure and Bonding – Foundations of Organic Chemistry

Chapter 1: Electronic Structure and Bonding

What is Organic Chemistry?

Organic chemistry is the study of compounds primarily composed of carbon atoms. Historically, organic compounds were thought to originate only from living organisms and were believed to possess a 'vital force.' In contrast, inorganic compounds were derived from minerals and lacked this force. This distinction was disproven when urea, an organic compound, was synthesized from inorganic ammonium cyanate.

• Organic compounds: Contain carbon; originally thought to come only from living organisms.

• Inorganic compounds: Typically do not contain carbon; derived from minerals.

• Example: Synthesis of urea from ammonium cyanate demonstrates that organic compounds can be made from inorganic sources.

What Makes Carbon So Special?

Carbon is unique among elements because of its position in the second row of the periodic table. It has the ability to form stable covalent bonds by sharing electrons, unlike elements to its left (which tend to lose electrons) or right (which tend to gain electrons).

• Electron sharing: Carbon forms four covalent bonds by sharing electrons.

• Versatility: This allows carbon to form a vast array of stable organic molecules.

• Comparison: Elements to the left of carbon (e.g., Be, B) tend to give up electrons; elements to the right (e.g., N, O, F) tend to accept electrons.

How Many Bonds Does an Atom Form?

The number of covalent bonds an atom forms is determined by its need to achieve a stable electron configuration (octet rule for most, duet for hydrogen).

• Oxygen: Forms 2 covalent bonds (e.g., in water, H2O).

• Nitrogen: Forms 3 covalent bonds (e.g., in ammonia, NH3).

• Carbon: Forms 4 covalent bonds (e.g., in methane, CH4).

• Hydrogen: Forms 1 covalent bond.

• Octet rule: Atoms (except hydrogen) are most stable when surrounded by 8 electrons.

Nonpolar and Polar Covalent Bonds

Covalent bonds can be classified based on the difference in electronegativity between the bonded atoms.

• Nonpolar covalent bond: Atoms have the same or similar electronegativity; electrons are shared equally (e.g., H–H, F–F, C–C, C–H).

• Polar covalent bond: Atoms have different electronegativities; electrons are shared unequally, resulting in partial charges (e.g., H–Cl, H–O, H–N).

• Electronegativity: The ability of an atom to attract electrons in a bond.

• Partial charges: Denoted as δ+ (partial positive) and δ– (partial negative).

The Effect of Electronegativity Difference on Bond Polarity

The polarity of a bond increases with the difference in electronegativity between the two atoms.

• Nonpolar covalent bond: Electronegativity difference < 0.5

• Polar covalent bond: Electronegativity difference between 0.5 and 1.9

• Ionic bond: Electronegativity difference > 1.9; electrons are not shared but transferred

• Bond polarity: Results in a dipole, with one end of the bond being more negative than the other.

Dipole Moment

The dipole moment is a quantitative measure of the polarity of a bond, defined as the product of the size of the charge and the distance between the charges.

• Formula: $\text{Dipole moment} = \text{charge} \times \text{distance}$

• Units: Debye (D)

• Relationship: The greater the difference in electronegativity, the greater the dipole moment and the more polar the bond.

Table: Dipole Moments of Some Common Bonds

Bond	Dipole moment (D)	Bond	Dipole moment (D)
H–C	0.4	C–C	0
H–N	1.3	C–N	0.2
H–O	1.5	C–O	0.7
H–F	1.7	C–F	1.6
H–Cl	1.1	C–Cl	1.5
H–Br	0.8	C–Br	1.4
H–I	0.4	C–I	1.2

• Interpretation: Bonds with higher dipole moments are more polar.

Additional info: These foundational concepts are essential for understanding the structure, reactivity, and properties of organic molecules. Mastery of electronic structure and bonding principles is critical for success in organic chemistry.