(c) Which is generally stronger, a s bond or a p bond? Explain.

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Understand the nature of sigma (σ) and pi (π) bonds: A sigma bond is formed by the head-on overlap of atomic orbitals, while a pi bond is formed by the side-to-side overlap of p orbitals.

Consider the overlap efficiency: Sigma bonds have greater overlap between orbitals compared to pi bonds, leading to a stronger bond.

Analyze the bond strength: The greater overlap in sigma bonds results in a stronger bond compared to pi bonds, which have less overlap and are generally weaker.

Think about bond formation: Sigma bonds are typically the first bonds formed between two atoms, while pi bonds are additional bonds that can form in double or triple bonds.

Conclude the comparison: Generally, a sigma bond is stronger than a pi bond due to the more effective overlap of orbitals in sigma bonds.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Bond Strength

Bond strength refers to the energy required to break a bond between two atoms. It is influenced by factors such as the type of orbitals involved in bond formation, the overlap of these orbitals, and the electronegativity of the atoms. Generally, stronger bonds have higher bond dissociation energies.

s and p Orbitals

s and p orbitals are types of atomic orbitals that describe the regions in an atom where electrons are likely to be found. s orbitals are spherical and can hold a maximum of two electrons, while p orbitals are dumbbell-shaped and can hold up to six electrons across three orientations. The shape and orientation of these orbitals affect how they overlap during bond formation.

Orbital Overlap

Orbital overlap is a key concept in chemical bonding that describes how atomic orbitals combine to form molecular bonds. The greater the overlap between orbitals, the stronger the bond formed. s bonds, formed from the head-on overlap of s orbitals, typically exhibit stronger overlap compared to p bonds, which involve side-to-side overlap of p orbitals, leading to generally weaker bonds.

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Textbook Question

(a) Which geometry and central atom hybridization would you expect in the series BH₄^-, CH₄, NH₄⁺?

Textbook Question

(b) What would you expect for the magnitude and direction of the bond dipoles in this series?

Textbook Question

(a) Draw a picture showing how two p orbitals on two different atoms can be combined to make a σ bond. (b) Sketch a π bond that is constructed from p orbitals.

Textbook Question

(d) Can two s orbitals combine to form a p bond? Explain.

Textbook Question

(b) Imagine that you could hold two atoms that are bonded together, twist them, and not change the bond length. Would it be easier to twist (rotate) around a single s bond or around a double 1s plus p2 bond, or would they be the same?