Study Guide for Covalent Compounds, Formulas, and Structures

Formation of Covalent Bonds:
- Covalent bonding occurs through the sharing of electrons between two or more atoms, allowing them to attain a noble-gas electronic configuration.
- Unlike ionic bonding, which involves the transfer of electrons and results in the attraction between oppositely charged ions, covalent compounds result from the physical attachment of atoms creating molecules.
Molecules:
- Compounds formed by covalently bonded groups of atoms are termed molecules.
- Lewis Structures:
- Chemists use Lewis electron-dot structures to represent the sharing of electrons in covalent bonds, displaying the outermost s and p electrons as dots around the atomic symbol.
- Noble Gas Configurations:
- Noble gases have full outer shells consisting of 2 electrons in the s electron shell and 6 in the p sublevel, totaling 8 electrons (the octet).
- The octet rule states that atoms strive to have 8 electrons in their outer shell (with hydrogen being an exception, achieving stability with 2).

Constructing Lewis Structures:
- For larger molecules, the initial task is to determine the “skeleton” (the arrangement of atoms).
- General guidelines for determining central atoms:
1. Carbon is usually central.
2. Hydrogen can only form one covalent bond and is never central.
3. Halogens generally are peripheral unless oxygen is absent.
4. Oxygen typically forms two covalent bonds and is rarely central.
- The construction of a Lewis structure involves several steps:
1. Count total valence electrons.
2. Adjust for charge if dealing with polyatomic ions.
3. Draw bonds between bonded atoms using pairs of electrons.
4. Distribute remaining electrons to complete outer octets.
5. If excess electrons remain, assign them to the central atom (though exceptions apply, especially with boron).
Examples of Lewis Structures:
- Methane (CH4) and Boron Trifluoride (BF3) demonstrate different central atom behaviors:
- Methane has a complete octet around carbon, while boron often forms compounds with incomplete octets.

Covalent Bonds:
- Double and triple bonds occur when two or three pairs of electrons are shared between atoms.
- An example is Sulfur Dioxide (SO2), which contains double bonds to meet the octet requirement.
Resonance Structures:
- Some molecules can be represented by multiple, valid Lewis structures, called resonance structures, each differing in the arrangement of double bonds.

Calculating Formal Charges:
- The formal charge formula determines the validity of a Lewis structure:
  $ext{Formal Charge} = ( ext{Valence Electrons}) - ( ext{Nonbonding Electrons}) - rac{1}{2} imes( ext{Bonding Electrons})$
- A structure with formal charges close to zero is generally preferred, and charges that add up to the overall charge of the ion are also considered.

Electronegativity:
- Developed by Linus Pauling, it quantifies an atom’s ability to attract electrons in a bond. It increases across periods (left to right) and decreases down groups (top to bottom).
- The difference in electronegativity values (99 EN) can determine bond polarity:
- Nonpolar covalent bonds (ΔEN = 0)
- Polar covalent bonds (ΔEN < 1.7)
- Ionic bonds (ΔEN ≥ 1.7)
Dipole Moment:
- The dipole moment quantifies the polarity of a bond, calculated as:
  $ext{Dipole Moment} = q imes r$
- The unit of measurement is the debye (1 D = 3.34 × 10^{-30} C·m).

Use VSEPR Theory to predict the three-dimensional arrangement based on the number of bonding electron pairs and nonbonding pairs around a central atom.
Various Geometries include:
- Linear (AX): 180°
- Planar triangle (AX2): 120°
- Tetrahedral (AX4): 109.5°
- Trigonal bipyramidal (AX5): 120° and 90°
- Octahedral (AX6): 90°
Molecular Polarity Rules:
1. A symmetrical molecule is likely nonpolar.
2. An asymmetrical molecule with polar bonds is likely polar.

Valence Bond (VB) Theory:
- Describes covalent bonds formed via the overlap of atomic orbitals, with specific arrangements (s, p) explaining molecule shapes.
Hybrid Orbitals:
- Hybrid orbitals result from mixing atomic orbitals to form new orbitals suited to bonding.
- sp3: (e.g., CH4) - tetrahedral arrangement (4 bond domains)
- sp2: (e.g., CH2O) - trigonal planar arrangement (3 bond domains + 1 lone pair)
- sp: (e.g., CO2) - linear arrangement (2 bond domains + 2 pi bonds)