Chapter 5 - Chemical Bonding I: The Lewis Model

Electron sea model 

It explains metallic bonding, where valence electrons are delocalized and free to move, creating a "sea" that bonds with positive metal ions, giving metals properties like conductivity and malleability.

Electronegativity

is the ability of an atom to attract shared electrons in a chemical bond.

How does electronegativity trend across the periodic table?

Electronegativity increases across a period from left to right and decreases down a group.

How do you determine bond type using electronegativity difference?

1. Find the Electronegativity Values:

   • Look up the electronegativity values of the two atoms involved in the bond.

2. Calculate the Difference:

   • Subtract the smaller electronegativity value from the larger one.

3. Interpret the Difference:

   • < 0.4: Nonpolar covalent bond (equal sharing of electrons)

   • 0.4 – 1.7: Polar covalent bond (unequal sharing of electrons)

   • > 1.7: Ionic bond (complete transfer of electrons)

What is molecular geometry's role in molecular polarity?

Symmetrical molecular geometries can cancel dipoles, making a molecule nonpolar, while asymmetrical shapes lead to polar molecules.

What is dipole moment and how is it affected by electronegativity?

is a measure of the separation of positive and negative charges in a molecule, representing the molecule’s polarity. It occurs when two atoms in a bond have different electronegativities, causing an unequal sharing of electrons. The larger the difference in electronegativity between the atoms, the larger the dipole moment.

What does percent ionic character indicate in a bond?

refers to how much the bond behaves like an ionic bond compared to a purely covalent bond.  The larger the electronegativity difference, the greater the percent ionic character.

What are resonance structures?

Different Lewis structures that represent the same molecule, differing only in electron positions, not atom positions. (Usually happens when there are two oxygens)

How is formal charge calculated? (practice this)

Formal Charge = # of Valence Electrons - (Nonbonding Electrons aka. lone pairs) + 1/2(Bonding Electrons aka double this because of the other atom it is bonded to)).

How does atomic radius affect bond length?

Larger atomic radii result in longer bonds; smaller atomic radii result in shorter bonds.

Order bond types by bond length and bond energy.

Single bonds: longest and weakest, double bonds: shorter and stronger, triple bonds: shortest and strongest.

Difference between VSEPR shapes (electron geometry) and Molecular geometry

The first considers lone pairs as bonds while the second one does not consider them as bond but just lone pairs

Name the VSEPR geometry and bond angle for a molecule with two electron domains.

Linear geometry, 180° bond angle.

What is the electron geometry and bond angle of a trigonal planar molecule?

Trigonal planar geometry, 120° bond angle.

Describe the bond angle and geometry of a molecule with four electron domains

Tetrahedral geometry with a 109.5° bond angle.

Describe the bond angle and geometry of a molecule with five electron domains

Trigonal Bipyramidal ( - 90° (axial), 120° (equatorial)

Describe the bond angle and geometry of a molecule with six electron domains

Octahedral - 90°

What are the possible molecular geometries for a trigonal bipyramidal electron geometry, based on 1, 2, and 3 lone pairs?

seesaw, T-shaped, and linear,

What are the possible molecular geometries for a Trigonal Planar electron geometry, based on 1 lone pair

Bent

What are the possible molecular geometries for a Tetrahedral electron geometry, based on 1 and 2 lone pairs

Trigonal Pyramidal, Bent

What are the possible molecular geometries for a Octahedral electron geometry, based on 1 and 2 lone pairs

Square Pyramidal, Square Planar

How do lone pairs affect bond angles?

They cause bond angles to decrease from their ideal values due to increased repulsion.

How do you predict molecular polarity?

Draw the Lewis Structure: Determine the arrangement of atoms and lone pairs.

1. Determine Bond Polarities: Assess the electronegativity differences to identify polar bonds.

2. Analyze Molecular Geometry: Use VSEPR theory to understand the shape of the molecule.

3. Evaluate Net Dipole Moment: Check if the dipole moments cancel out or not:

   • Polar Molecule: If dipole moments do not cancel (net dipole moment ≠ 0).

   • Nonpolar Molecule: If dipole moments cancel (net dipole moment = 0).

What is dual polarity in a molecule?

It occurs when a molecule has both polar and nonpolar regions, often due to different functional groups.

Nonpolar Covalent Bonds

Bonding eletrons shared equally between atoms

Polar Covalent Bond

Bonding electrons shared unequally between two atoms (usually due to

How does molecular shape affect the dipole moment?

Symmetrical shapes can cancel out dipole moments, resulting in nonpolar molecules, while asymmetrical shapes often lead to polar molecules.

What does the term "bond dissociation energy" refer to?

is the energy required to break a bond, indicating bond strength; higher energy means a stronger bond.

How do expanded octets affect molecular structure?

They allow molecules to accommodate more electrons, leading to structures like trigonal bipyramidal and octahedral shapes.

What are incomplete octets, and which elements commonly have them?

They occur when atoms have fewer than eight electrons in their valence shell. Elements like boron and beryllium often have them.

What is a free radical in terms of Lewis structures?

It is a molecule with an odd number of electrons, resulting in an unpaired electron, making it highly reactive.

How to determine a free radical?

a molecule with an odd number of total valence electrons often ends up with an unpaired electron, which makes it a free radical. This unpaired electron gives the molecule reactive properties, as it seeks to pair up with another electron to achieve stability.

Rules of Formal Charge

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

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

3. Small (or zero) formal charges on individual atoms are better than large ones.

4. When formal charge cannot be avoided, negative formal charge should reside on the most electro- negative atom.