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chemistry ch.1 pt.1

1. Bond Formation

Bond formation is the process by which atoms join together to form molecules or compounds. This usually occurs to achieve a more stable electron configuration, often by satisfying the octet rule (having eight valence electrons, with the exception of hydrogen and helium which aim for two). Energy is released when bonds form, indicating that the resulting molecule is at a lower, more stable energy state than the separated atoms.

2. Orbitals
2.1 Atomic Orbitals

Atomic orbitals are regions around the nucleus where electrons are most likely to be found. They are characterized by quantum numbers and have specific shapes (s, p, d, f).

2.2 Molecular Orbitals and Hybridization

When atoms bond, their atomic orbitals can combine to form molecular orbitals. This process often involves hybridization, where atomic orbitals mix to form new, degenerate hybrid orbitals that are suitable for bonding. Common hybridizations include:

  • sp^3: Forms four equivalent bonds (e.g., methane, \text{CH}_4)

  • sp^2: Forms three equivalent bonds and one p orbital for double bonding (e.g., ethene, \text{C}2\text{H}4)

  • sp: Forms two equivalent bonds and two p orbitals for triple bonding or two double bonds (e.g., ethyne, \text{C}2\text{H}2)

3. Formal Charge

Formal charge is the hypothetical charge an atom would have if all electrons in a bond were shared equally between the atoms, regardless of electronegativity. It helps in determining the most plausible Lewis structure for a molecule.

3.1 Calculation of Formal Charge

The formal charge of an atom in a molecule can be calculated using the formula:

\text{Formal Charge} = (\text{Valence electrons}) - (\text{Non-bonding electrons}) - \frac{1}{2}(\text{Bonding electrons})

3.2 Rules for Determining Most Stable Structures

  • The sum of formal charges in a neutral molecule must be zero.

  • The sum of formal charges in an ion must equal the charge of the ion.

  • Lewis structures with smaller formal charges (closer to zero) on individual atoms are generally more stable.

  • Negative formal charges should ideally reside on more electronegative atoms.

4. Bond Polarity

Bond polarity describes the unequal sharing of electrons between atoms in a covalent bond, leading to a separation of electric charge. It is determined by the difference in electronegativity between the bonded atoms.

4.1 Electronegativity

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. The greater the difference in electronegativity between two bonded atoms, the more polar the bond.

4.2 Types of Covalent Bonds

  • Nonpolar Covalent Bond: Occurs when two atoms share electrons equally or nearly equally, typically when the electronegativity difference is less than 0.4. (e.g., \text{H}2, \text{Cl}2)

  • Polar Covalent Bond: Occurs when there is an unequal sharing of electrons, resulting in partial positive (\delta+) and partial negative (\delta-) charges on the bonded atoms. The electronegativity difference is usually between 0.4 and 1.8-2.0. (e.g., \text{H-Cl}, \text{H-F})

4.3 Dipole Moments

A polar bond creates a bond dipole, represented by an arrow pointing toward the more electronegative atom. The overall polarity of a molecule (molecular dipole moment) depends on both the polarity of its individual bonds and the molecule's geometry.