Molecular Geometry and Bonding Theories

1. Electronegativity

1.1. Ability to attract electrons, Periodic trend, Increasing electronegativity across a period, and Decreasing electronegativity down a group.

1.2. The difference in electronegativity between atoms in a bond determines its polarity.

1.3. Negative charge can be designated as δ− partial positive charge can be designated as δ+

2. Bond Polarity

2.1. Electronegativity difference, Polar bonds, Non-polar bonds, and Dipole moment.

2.1.1. Example: Water molecule (H2O): Oxygen has a higher electronegativity than hydrogen, so the shared electrons are pulled closer to the oxygen atom, making the oxygen slightly negative and the hydrogen atoms slightly positive. This polarity is what allows water molecules to interact with each other through hydrogen bonds.

2.2. Nonpolar bonds: When two atoms have similar electronegativity, the electrons are shared equally, resulting in a nonpolar bond.

2.3. Polar bonds: When two atoms have significantly different electronegativities, the electrons are shared unevenly, creating a polar bond.

2.4. Ionic bonds: A very large difference in electronegativity can lead to an ionic bond, where one atom essentially "donates" its electron to the other, creating fully charged ions.

3. Resonance Structures

3.1. Multiple Lewis structures, Delocalized electrons, and Average representation of bonding.

3.1.1. Example: Nitrate ion (NO3-): This ion can be represented by three resonance structures where the double bond is located between the nitrogen atom and each oxygen atom in turn. The actual structure is a hybrid of these three resonance forms, meaning the bonds between N and O are equivalent and have a bond order of 1.33

4. Bond Enthalpy (Bond Strength):

4.1. A measure of the energy required to break a chemical bond.

4.2. A double bond is stronger than a single bond, and a triple bond is stronger than both a single and double bond between the same atoms.

5. Bonding Theories

5.1. Ionic Bonding

5.1.1. Transfer of electrons Metal and non-metal interaction Electrostatic attraction Cation and anion formation

5.1.1.1. Example: Sodium chloride (NaCl) - Sodium (Na) readily loses an electron to chlorine (Cl) forming Na+ and Cl- ions, which are attracted to each other.

5.2. Covalent Bonding

5.2.1. Sharing of electrons Between non-metals Single, double, and triple bonds Polar covalent bonds Non-polar covalent bonds

5.2.1.1. Example: Water (H2O) - Oxygen shares electrons with two hydrogen atoms, creating a polar covalent bond.

5.3. Metallic Bonding

5.3.1. Delocalized electrons Sea of electrons Conductivity Malleability and ductility

5.3.1.1. Exampl: copper, silver, gold, iron, aluminum, sodium, magnesium, and most other pure metals

6. VSEPR Model

6.1. Electron domains, Minimizing repulsion, Predicting molecular geometry, Electron pair geometry, and Lone pair effects.

7. Lewis Structures

7.1. Lewis structures are used to determine the electron domains around a central atom, which are then used to predict molecular geometry according to VSEPR theory.

7.2. Valence electrons representation, Octet rule (with exceptions), Single bonds, Double bonds, and Triple bonds.

8. Octet Rule

8.1. The tendency of atoms to prefer to have 8 electrons in the valence shell. When atoms have fewer than 8 electrons, they tend to react and form more stable compounds.

8.2. Exception to the rule: Certain elements like B and Be can have less than 8 electrons in their outermost shell, while elements like P and S can exceed the octet rule due to the availability of d orbitals.

9. Hybrid Orbitals and Molecular Orbitals

9.1. Hybrid orbitals are formed by combining atomic orbitals within the same atom

9.2. Molecular orbitals are formed by combining atomic orbitals from different atom