University Notes: Chemistry Chapter 4 - Covalent Bonding and Molecular Compounds

General Theme: This chapter explores how atoms share electrons to form molecules and the subsequent impact of these substances on human health and the environment.
Impact of Chemical Interactions: Conditions such as depression, arthritis, and carbon monoxide poisoning are all dictated by chemical interactions resulting from atoms forming molecules.
Stated Learning Objectives:
- Use Lewis structures to explain the structure of molecules.
- Write and name molecular compounds.
- Justify the phrase "I cannot exist without molecules."
- Explain the relationship between the number of atoms present in a molecular compound and their mass.
- Contrast the characteristics of polar and nonpolar covalent molecules.

Case Studies in Molecular Impact:
- Emotional and Mental Health (Deborah): Deborah experienced severe anxiety and hopelessness. Her physician prescribed an antidepressant, which utilized chemical interactions to help her feel calmer and resolve her difficulties.
- Physical Ailments (Andrea): Reaching middle age, Andrea suffered from stiff, painful joints in the morning due to arthritis. Aspirin, a molecular compound, provided relief and allowed her to function despite the pain.
- Lethal Hazards (Carbon Monoxide): In $2008$ , $13$ people (including $4$ Fire Department personnel) in Seattle were hospitalized for dizziness and headaches caused by carbon monoxide poisoning from a running van. It is noted that approximately $3000$ people die from carbon monoxide poisoning annually.
Commonality: Both the benefits (relief from depression and arthritis) and the dangers (poisoning) originate from compounds formed through covalent bonding.

Distinction from Ionic Bonds:
- Ionic Bonds: Investigated in Chapter 3, these form when metal atoms transfer electrons to nonmetal atoms so both achieve a noble gas electron configuration.
- Covalent Bonds: Form when nonmetals share electrons with other nonmetals to fill their valence shells.
Behavior of Nonmetals:
- Nonmetals are located in the top right corner of the periodic table.
- They rarely transfer electrons to other atoms; instead, they seek to fill their valence shells by either taking electrons from a metal (ionic) or sharing with another nonmetal (covalent).
Definitions and Terminology:
- Molecule: The resulting group of atoms when nonmetals form a covalent bond.
- Molecular Compounds: These are identical to covalent compounds. The textbook uses these terms interchangeably as they both refer to substances composed of molecules.
- Molecule Scale: The molecule is considered the smallest particle of a covalent compound.

Gilbert Newton Lewis ( $1875$ - $1946$ ):
- Lewis is credited with making chemical bonding comprehensible through dot structures.
- The Cube Model: Inspired by Richard Abegg's "octet rule" (atoms are most stable with $8$ electrons in their outer shell), Lewis originally visualized electron positions as concentric cubes. The dots at the eight corners of these cubes represented electron locations.
- Origin of Covalent Bonding: Lewis realized atoms could complete these cubes (valence shells) by sharing electrons. He published this in his classic $1916$ paper, The Atom and the Molecule.
- Linus Pauling: Later built upon Lewis's work by calculating electronegativity values to predict bond polarity.
Biographical and Professional Details:
- Lewis was a faculty member at MIT before moving to the University of California-Berkeley in $1912$ as Dean of the College of Chemistry.
- At Berkeley, he revolutionized chemistry education by emphasizing original thinking and concepts rather than rote memorization.
- Scientific Contributions: Investigated chemical thermodynamics, isolated "heavy water" (containing a hydrogen isotope), and coined the term "photon" for the smallest unit of radiant energy.
- Nobel Prize: Despite being nominated $35$ times, Lewis never won the Nobel Prize, though he remains one of the most distinguished chemists in history.

The "Tug-of-War" Metaphor: Atoms in a covalent bond do not always share electrons equally. It is described as a tug-of-war between well-matched teams where both sides try to gain the electrons without giving up their own.
Electronegativity:
- Definition: A measure of an atom's ability to draw electrons from other atoms in a molecule toward itself.
- Periodic Trends: Fluorine ( $F$ ) is the most electronegative element (upper right corner). Electronegativity decreases diagonally down and to the left toward Cesium ( $Cs$ ), which is the least electronegative (Fig. 4.2).
- Mechanism of Sharing: Electron sharing occurs when both atoms have relatively high electronegativity (close to Fluorine). The atom closer to Fluorine exerts a stronger pull, but cannot fully remove the electrons from the other atom.

Types of Covalent Bonds:
- Single Bonds: Two atoms share one pair of electrons ( $2$ electrons total). Represented by one line (e.g., $H-O-H$ in Water, $H_2O$ ).
- Double Bonds: Two atoms share two pairs of electrons ( $4$ electrons total). Represented by two lines (e.g., $O=C=O$ in Carbon Dioxide, $CO_2$ ).
- Triple Bonds: Two atoms share three pairs of electrons ( $6$ electrons total). Represented by three lines (e.g., $N ≡ N$ in Nitrogen, $N_2$ ).
Bond Energy:
- Definition: The energy required to break a specific chemical bond.
- Physics of Attraction: Electrons (negative) attract protons (positive) in the nuclei of both atoms. More shared electron pairs lead to stronger attraction and higher bond energy.
- Measurement Unit (The Joule): Expressed in kilojoules per mole ( $kJ/mol$ ). A Joule is the energy required to move a $1\,kg$ mass $1\,m$ at an acceleration of $1\,m/s^2$ .
Bond Length:
- Definition: The distance between the nuclei of covalently bonded atoms.
- The Spring Metaphor: A bond is like a spring that vibrates around an equilibrium length. If compressed, repulsive forces lengthen it; if stretched, attractive forces shorten it (Fig. 4.3).
- Equilibrium: The bond length is the point where attractive and repulsive forces (between protons in the nuclei) are balanced.

Stability Trends: Shorter bond lengths correlate with greater stability because the bond energy is stronger, making the bond harder to break.
- Triple Bonds: Strongest energy, shortest length, highest stability.
- Double Bonds: Greater energy and shorter length than single bonds.
- Single Bonds: Lowest energy and longest length.
Data from Table 4.1 (Bond Comparison):
- C-C (Single): Bond Energy: $348\,kJ/mol$ ; Length: $154\,pm$ .
- C=C (Double): Bond Energy: $614\,kJ/mol$ ; Length: $134\,pm$ .
- C≡C (Triple): Bond Energy: $839\,kJ/mol$ ; Length: $120\,pm$ .
- C-O (Single): Bond Energy: $360\,kJ/mol$ ; Length: $143\,pm$ .
- C=O (Double): Bond Energy: $802\,kJ/mol$ ; Length: $120\,pm$ .
- C≡O (Triple): Bond Energy: $1077\,kJ/mol$ ; Length: $113\,pm$ .
- N≡N (Triple): Bond Energy: $945.4\,kJ/mol$ ; Length: $110\,pm$ .
- O=O (Double): Bond Energy: $498.3\,kJ/mol$ ; Length: $121\,pm$ .