CHEM 142_Lecture_4

CHEM 142: The Global Impact of Chemistry for Engineering

Instructor: Dr. Levent Inci

Copyright © 2020 W. W. Norton & Company

Chapter 9: Molecular Geometry: Shape Determines Function

Authors: Gilbert Kirss, Bretz, Foster

Copyright © 2020 W. W. Norton & Company

Chapter Outline

• 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR)

• 9.3 Polar Bonds and Polar Molecules

Learning Outcomes

• Predict bond angles and shapes of molecules utilizing VSEPR theory and the concept of steric number.

• Determine substance polarity, identifying whether a molecule is polar or nonpolar by analyzing molecular structure and the types of chemical bonds present.

Biological Activity and Molecular Shape

Molecular shape plays a critical role in influencing both chemical and physical properties of substances, affecting the interactions and reactivity of molecules. For instance, in biological systems, the three-dimensional structure of a molecule is vital for its biological function; enzyme-substrate interactions depend fundamentally on the shape specificity of active sites. The compatibility of shapes among biomolecules enables selective interactions that are essential for biological processes such as signal transduction and metabolic pathways.

Molecular Shape

Lewis Structures:Lewis structures provide a two-dimensional representation of atoms and covalent bonds, illustrating how different atoms are connected. However, they fall short in conveying the three-dimensional spatial orientation that is crucial for comprehending molecular properties and reactivity. Thus, while useful for visualizing connectivity, they do not fully represent the complexities of molecular shapes.

Molecular Models:Three-dimensional molecular models are invaluable for displaying the orientations and angles of chemical bonds more effectively. These models enhance understanding of physicochemical properties such as polarity, reactivity, and are also crucial in predicting molecular behavior during chemical reactions.

Bond Angle

• Definition: The bond angle is the angle in degrees formed by the bond lines connecting the centroids of three covalently bonded atoms. The bond angles can significantly affect a molecule's reactivity, stability, and interactions with other molecules.

Valence-Shell Electron-Pair Repulsion Theory (VSEPR)

VSEPR theory provides a systematic approach to predicting the spatial arrangement of atoms within a molecule, aiming to minimize electron pair repulsions associated with the bonds and lone pairs of electrons.

• Electron-Pair Geometry: This refers to the three-dimensional arrangement of bonding pairs and lone pairs of electrons around a central atom. Understanding electron pair geometry is essential for predicting molecular reactivity and properties.

• Molecular Geometry: This pertains specifically to the three-dimensional arrangement of only the atoms within a molecule, which can vary from the electron-pair geometry, especially in cases with lone pairs.

Determining Electron-Pair Geometry

To determine the geometry of a molecule, follow these steps:

1. Draw the Lewis structure for the molecule or ion, incorporating all valence electrons to ensure electron accounting is accurate.

2. Calculate the Steric Number (SN):

   • SN = (number of atoms bonded to the central atom) + (number of lone pairs on the central atom).

3. Determine the optimal spatial arrangement to minimize electron electron repulsion, based on the calculated steric number.

Molecular Geometry Without Lone Pairs

• SN = 2: Linear geometry (two atoms bonded) with bond angle of 180°.

• SN = 3: Trigonal planar geometry (three atoms bonded) with bond angle of 120°.

• SN = 4: Tetrahedral geometry (four atoms bonded) with bond angle of 109.5°.

• SN = 5: Trigonal bipyramidal geometry (five atoms bonded) featuring bond angles of 90° and 120°.

• SN = 6: Octahedral geometry (six atoms bonded) with a bond angle of 90°.

Central Atoms with Lone Pairs

The presence of lone pairs significantly affects molecular geometry, often resulting in deviations from the ideal bond angles due to greater repulsion exerted by lone pairs compared to bonding pairs.Example: In sulfur dioxide (SO2) with an SN = 3, the structure features two bonding pairs and one lone pair, yielding a bent molecular shape rather than a straight one.

Bond Angles with Lone Pairs

Bond angles in molecular structures containing lone pairs can often be reduced compared to the expected ideal geometries due to increased repulsion from lone pairs. The presence of additional lone pairs generally enhances this electron repulsion, leading to smaller bond angles than predicted.

Bond Polarity Review

Types of Bonds:

• Nonpolar covalent bonds: Electrons are shared equally between atoms, leading to a balanced charge distribution.

• Polar covalent bonds: Electrons are shared unequally, resulting in a dipole moment due to charge separation.

• Ionic bonds: Electrons are transferred from one atom to another, creating charged particles via the formation of cations and anions.

Polar Bonds and Polar Molecules

Requirements for Polar Molecules:

• Must exhibit polar bonds.

• Have a non-uniform distribution of these polar bonds that creates an overall molecular dipole moment, affecting physical properties such as solubility and boiling point.

Bond Dipole and Polar Molecules:

• Definitions:

  • Bond dipole: A quantitative measure of charge separation exhibited between two atoms in a bond, arising from differences in electronegativity.

  • Polar molecule: A molecule where the sum of its bond dipoles does not cancel out and results in a net dipole moment, making it polar.

Measuring Polarity

• Dipole moment (μ): This measurement quantifies the overall polarity of a molecule and is expressed in debyes (D), where 1 D = 3.34 × 10-30 C·m.

• Dipole Moments of Polar Molecules:

  • HF: Permanent dipole of 1.91 D

  • H2O: Dipole moment of 1.85 D

Practice: Predicting the Polarity of Formaldehyde (CH2O)

1. Determine the presence of a permanent dipole by analyzing the symmetry and bond orientations of the molecule.

2. Evaluate electronegativities: H (2.1), C (2.5), and O (3.5); it can be noted that there is a dipole directed toward the oxygen atom, indicating the presence of a permanent dipole moment.

Chapter 10: Intermolecular Forces: The Uniqueness of Water

Authors: Gilbert Kirrs, Bretz, Foster

Copyright © 2020 W. W. Norton & Company

Chapter Outline

• 10.1 Intramolecular Forces versus Intermolecular Forces

• 10.2 Dispersion Forces

• 10.3 Interactions Involving Polar Molecules

Learning Outcomes

• Explain the origin and understanding characteristics of various intermolecular forces, including dispersion forces, ion-dipole interactions, dipole-dipole interactions, and hydrogen bonds.

Intramolecular vs. Intermolecular Forces

• Intramolecular forces: These are the attractive and repulsive forces operating within a molecule that hold the atoms together, such as covalent bonds.

• Intermolecular forces: Forces that act between separate molecules, crucially influencing physical properties such as boiling points, melting points, and solubility of substances.

States of Matter

• Solid: Has a defined shape and volume, characterized by strong interatomic forces leading to a fixed structure.

• Liquid: Maintains a defined volume but lacks a definite shape; intermolecular forces allow the molecules to flow past one another, resulting in fluidity.

• Gas: Does not have a defined shape or volume; intermolecular forces are minimal, allowing gas molecules to expand and occupy available space.

Boiling Points and Atomic Number

Boiling points tend to increase with the atomic number of elements due to stronger intermolecular attractions requiring more energy to separate, resulting in higher boiling points for substances with larger atomic masses.

Interactions of Nonpolar Molecules

Dispersion (London) Forces:These are the weakest intermolecular forces present between nonpolar molecules, arising from temporary dipoles created by momentary fluctuations in electron distribution within atoms. This results in temporary charge imbalances that can affect neighboring molecules.

Polarizability

Polarizability refers to the tendency of an electron cloud around a molecule to become distorted, leading to the generation of a temporary dipole. Larger atoms or molecules, having more electrons, exhibit higher polarizability and consequently, stronger dispersion forces.

Strength of Dispersion Forces

Factors Affecting Strength:

• Size of atoms or molecules: Larger atoms have more electrons, leading to increased polarizability and stronger dispersion forces.

• Shape of molecules: Molecules with a more extended surface area experience stronger dispersion forces due to greater contact and interaction with neighboring molecules.

Effect of Size on Dispersion

Example Table: Boiling points increase with the molecular mass of halogens:

• F2 (85 K) < Cl2 (239 K) < Br2 (332 K) < At2 (610 K)

Effect of Shape on Dispersion

Molecules with a more spread-out conformation possess stronger dispersion forces and subsequently, exhibit higher boiling points when compared to more compact molecular structures, owing to greater surface interactions.

Interactions Involving Polar Molecules

Ion-Dipole Interaction:This interaction occurs between charged ions and polar molecules and is fundamental for understanding solubility and reaction kinetics, particularly in electrochemical processes. Ion-dipole interactions play a crucial role in the solvation of ions in solution.

Sphere of Hydration

The sphere of hydration refers to the cluster of water molecules that surround and stabilize an ion in solution. This reduces the ion's charge density, which further enhances its solubility due to the effective interaction between the polar water molecules and the ion.

Dipole-Dipole Interactions

These interactions are attractive forces that occur between the positive end of one polar molecule and the negative end of another. They are essential in defining the boiling points and solubility of polar substances, having significant implications for the properties of liquids and solutions.

Hydrogen Bond

This is a notably strong type of dipole-dipole interaction that occurs when hydrogen is covalently bonded to electronegative elements like nitrogen (N), oxygen (O), or fluorine (F). Hydrogen bonds greatly influence the physical properties of compounds such as water, resulting in its unique characteristics.

Dipole-Induced Dipole

This phenomenon occurs when a polar molecule induces a dipole in a neighboring nonpolar molecule due to their close proximity. This interaction can affect properties like solubility and reactivity of nonpolar substances when mixed with polar solvents.

Relative Strengths of Intermolecular Forces

Example Table: Illustrates the comparative effects of various intermolecular forces on solubility and states of matter:

• NaCl dissolves in water demonstrating ion-dipole interactions, which are crucial for ionic compounds.

• Water expands when it freezes due to hydrogen bonding, which causes a significant disruption in close packing of molecules.

• Formaldehyde (CH2O) has a notably higher boiling point compared to ethane (C2H6) due to stronger dipole-dipole interactions, as ethane primarily exhibits weaker dispersion forces.