Gen Chem - Ch. 6 and 7

What are the three learning objectives for Chapter 06: Lewis Structures and Bonding?

1. Understand Lewis Dot Structures and Geometries.

2. Draw dipole moments and identify polar and non-polar molecules.

3. Describe intermolecular forces between like and unlike molecules.

What is a Dipole Moment?

In heteroatomic molecules, electrons are not shared equally in covalent bonds. Electrons are more likely to be near one nucleus, resulting in partial positive (δ+) and negative (δ-) charges on the atoms. This separation of charge is a dipole moment.

Example: In HF, the bond is polar covalent with H being δ+ and F being δ-.

What is Electronegativity?

The characteristic ability of an element to attract electrons in a covalent bond. It is related to ionization energy and electron affinity. The more electronegative atom in a bond will have a partial negative charge. It increases from left to right and decreases from top to bottom on the periodic table.

What are the trends for Electronegativity?

Electronegativity increases across a period (left to right) and decreases down a group (top to bottom). Fluorine (F) is the most electronegative element.

What is a Polar Molecule?

A molecule that has a net dipole moment, with distinct negative and positive regions. This occurs when the vector sum of the individual bond dipoles is not zero. Polar molecules attract each other; the positive end of one molecule is attracted to the negative end of another.

How is the overall polarity of a molecule predicted?

By taking the vector sum of all the bond dipoles. Bond dipoles are usually shown as crossed arrows, with the arrowhead pointing toward the negative end.

When is a molecule typically NONPOLAR?

When it is symmetrical and all bond dipoles cancel out. All five basic molecular shapes (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral) are symmetrical and nonpolar if all atoms attached to the central atom are identical.

When is a molecule typically POLAR?

1. If all atoms attached to the central atom are NOT the same. OR

2. If there are one or more lone pairs on the central atom (with some exceptions). Example: Water (H₂O) and ammonia (NH₃) are polar due to their lone pairs and non-canceling bond dipoles.

What are the three learning objectives for Chapter 07: Bonding Theories?

1. Predict structure based on VSEPR and valence bond theory.

2. Utilize molecular orbital theory to predict electronic structure of molecules.

3. Understand the advantages and disadvantages of each theory and use them appropriately.

What is Valence Bond (VB) Theory?

A theory that proposes a chemical bond is formed when atomic orbitals overlap, and a pair of electrons with paired (opposite) spins is shared in the overlapping orbitals. The extent of overlap is related to bond strength.

Example: In H₂, the 1s orbitals from each H atom overlap to form a covalent bond.

What is a key difficulty with the basic Valence Bond Theory?

It often fails to predict correct molecular geometries.

Example: Basic VB theory predicts CH₂ with 90° bond angles (based on carbon's p orbitals), but methane (CH₄) has four equivalent bonds with 109.5° bond angles.

What are Hybrid Orbitals in Valence Bond Theory?

Orbitals formed by mixing (hybridizing) two or more atomic orbitals on the same atom. They are used to explain molecular geometries that pure atomic orbitals cannot.

What are the properties of hybrid orbitals?

1. They are blended orbitals from the hybridization process.

2. The number of hybrid orbitals formed equals the number of atomic orbitals mixed.

3. The sum of the exponents in the hybrid orbital notation (e.g., sp³) must equal the number of atomic orbitals used.

Summarize the common hybrid orbital schemes and their geometries.

sp (1 s + 1 p): Linear, 180°
sp² (1 s + 2 p): Trigonal Planar, 120°
sp³ (1 s + 3 p): Tetrahedral, 109.5°
sp³d (1 s + 3 p + 1 d): Trigonal Bipyramidal, 90° & 120°
sp³d² (1 s + 3 p + 2 d): Octahedral, 90°

What is the hybridization of oxygen in OCl₂?

The central O atom has 2 bonding domains and 2 lone pairs, for a total of 4 electron regions. This corresponds to sp³ hybridization.

What is a Sigma (σ) Bond?

A covalent bond formed by the end-to-end overlap of atomic orbitals along the internuclear axis. It is the first bond formed between any two atoms. All single bonds are sigma bonds. Can be formed by s-s, s-p, p-p (end-on), or hybrid orbital overlap.

What is a Pi (π) Bond?

A covalent bond formed by the side-by-side overlap of p orbitals, creating electron density above and below the internuclear axis. It is the second (and third) bond in a multiple bond (double or triple bond).

How many σ and π bonds are in a double bond? In a triple bond?

A double bond contains 1 σ bond and 1 π bond. A triple bond contains 1 σ bond and 2 π bonds.

What is Molecular Orbital (MO) Theory?

A theory where molecular orbitals are formed from the linear combination of atomic orbital (LCAO) wave functions. These molecular orbitals are delocalized over the entire molecule.

What is a Bonding Molecular Orbital?

Formed by the constructive interference (addition) of atomic wave functions. Electron density builds up between the nuclei, lowering the energy of the molecule and making it more stable than the separate atoms. Denoted as σ or π.

What is an Antibonding Molecular Orbital?

Formed by the destructive interference (subtraction) of atomic wave functions. Creates a nodal plane between the nuclei, reducing electron density between them. It is higher in energy than the parent atomic orbitals and destabilizes the molecule. Denoted as σ* or π*.

What are the rules for filling Molecular Orbitals with electrons?

1. Electrons fill the lowest-energy orbitals available (Aufbau Principle).

2. No more than two electrons, with paired spins, can occupy any orbital (Pauli Exclusion Principle).

3. Electrons spread out with unpaired spins over orbitals of the same energy before pairing up (Hund's Rule).

How is Bond Order calculated in MO theory?

Bond Order = (number of bonding electrons - number of antibonding electrons) / 2

What does the Bond Order tell you?

A positive bond order indicates a stable molecule. Bond order correlates with bond strength and bond length: higher bond order = stronger, shorter bond.

Using MO theory, why does He₂ not form?

He has 2 valence electrons each, so He₂ has 4 total electrons. They fill both the σ₁ₐ (bonding) and σ₁ₐ* (antibonding) orbitals. Bond Order = (2 - 2)/2 = 0. A bond order of zero means the molecule is no more stable than the separate atoms.

What is the key difference in the MO energy level diagram for B₂ and C₂ compared to O₂ and F₂?

For B₂ and C₂ (atoms Li through N), the σ₂ₚ orbital is higher in energy than the π₂ₚ orbitals. For O₂ and F₂ (atoms O and higher), the σ₂ₚ orbital is lower in energy than the π₂ₚ orbitals.

How does MO theory successfully explain the paramagnetism of O₂?

The MO diagram for O₂ shows two unpaired electrons in the degenerate π₂ₚ* antibonding orbitals. Paramagnetism (attraction to a magnetic field) is caused by the presence of unpaired electrons, which Lewis structures and VB theory fail to predict.

Compare the strengths and weaknesses of Valence Bond (VB) Theory and Molecular Orbital (MO) Theory.

VB Theory:

Strengths: Intuitive, explains bond formation and geometry well, uses hybrid orbitals.

Weaknesses: Does not adequately explain resonance or paramagnetism in molecules like O₂.

MO Theory:

Strengths: Correctly predicts bond orders, magnetism, and delocalized electrons (resonance).

Weaknesses: Complex energy level diagrams, requires extensive calculations for larger molecules, less intuitive.