Covalent Bonding Study Guide

Chapter 16: Covalent Bonding\n\n# 16.1 The Nature of Covalent Bonding\n\n* Introduction to Covalent Bonds\n * Covalent bonding involves atoms sharing electrons to acquire the electron configurations of noble gases, often forming an octet ($8$ valence electrons).\n * This differs from ionic bonding where electrons are transferred; covalent bonding is like a \"tug of war\" that ends in a standoff.\n * Combinations of nonmetallic elements in Groups $4A$, $5A$, $6A$, and $7A$ are most likely to form covalent bonds.\n * Example: Hydrogen ($H_2$) consists of two hydrogen atoms sharing a pair of electrons. Each hydrogen achieves the electron configuration of helium ($2$ valence electrons).\n\n* Single Covalent Bonds\n * A single covalent bond is a bond in which two atoms share a pair of electrons.\n * Structural Formula: A chemical formula that shows the arrangement of atoms in molecules and polyatomic ions. Represented by dashes (e.g., $H-H$). Each dash indicates one pair of shared electrons.\n * Ionic compounds describe formula units (lowest whole-number ratio, like $CuO$ for Copper(II) Oxide), while molecular formulas describe the actual number of atoms in a molecule (e.g., $H_2$, $H_2O$).\n\n* Unshared Pairs (Lone Pairs)\n * Unshared pairs are pairs of valence electrons that are not shared between atoms, also known as lone pairs or nonbonding pairs.\n * Example: In Fluorine ($F_2$), each fluorine atom has three unshared pairs and one shared pair.\n\n* Single Covalent Bond Examples and Electron Dot Structures\n * Water ($H_2O$): Oxygen forms two single covalent bonds with two hydrogen atoms. Oxygen has two unshared pairs.\n * Ammonia ($NH_3$): Nitrogen forms three single covalent bonds with three hydrogen atoms. Nitrogen has one unshared pair.\n * Methane ($CH_4$): Carbon forms four single covalent bonds with four hydrogen atoms. Carbon has no unshared pairs.\n\n* The Case of Carbon's Four Bonds\n * Carbon's ground state is $1s^2 2s^2 2p^2$. Theoretically, it might only form two bonds. However, one $2s$ electron is promoted to a vacant $2p$ orbital ($1s^2 2s^1 2p^3$), providing four electrons for bonding.\n * The stability of $CH_4$ compensates for the small energy cost of this electron promotion.\n\n* Double and Triple Covalent Bonds\n * Double Covalent Bond: Involves two shared pairs of electrons ($4$ electrons total).\n * Triple Covalent Bond: Involves three shared pairs of electrons ($6$ electrons total).\n * Example: Nitrogen ($N_2$) forms a triple bond ($:N \equiv N:$) to achieve an octet. Oxygen ($O_2$) is an exception; though it has double-bond character, it has two unpaired electrons (paramagnetic).\n * Carbon dioxide ($CO_2$): Contains two carbon-oxygen double bonds ($O=C=O$).\n\n* Diatomic Elements Summary\n * Fluorine ($F_2$): Greenish-yellow reactive toxic gas; added to toothpaste and water.\n * Chlorine ($Cl_2$): Greenish-yellow reactive toxic gas; used in bleaching agents.\n * Bromine ($Br_2$): Dense red-brown liquid; used in photographic emulsions.\n * Iodine ($I_2$): Dense gray-black solid with purple vapors; antiseptic tincture.\n * Hydrogen ($H_2$): Colorless, odorless, tasteless gas; lightest element.\n * Nitrogen ($N_2$): Colorless, odorless, tasteless gas; $80\%$ of air by volume.\n * Oxygen ($O_2$): Colorless, odorless, tasteless gas vital for life; $20\%$ of air by volume.\n\n* Coordinate Covalent Bonds\n * Definition: A covalent bond in which one atom contributes both bonding electrons.\n * Represented by an arrow pointing from the donor to the receiver in structural formulas.\n * Example: Carbon Monoxide ($CO$). Carbon ($4$ valence electrons) and Oxygen ($6$ valence electrons) share two pairs initially. Oxygen then donates one of its lone pairs to Carbon so both achieve octets ($C ← O$).\n * Example: Ammonium Ion ($NH_4^+$). Forms when a hydrogen ion ($H^+$) is attracted to the unshared pair of an ammonia molecule ($NH_3$).\n\n* Bond Dissociation Energy\n * Definition: The total energy required to break the bond between two covalently bonded atoms.\n * High dissociation energy correlates with high chemical stability and low reactivity.\n * Specific Energies:\n * $H-H$ bond: $435\,kJ\,mol^{-1}$\n * $C-C$ single bond: $347\,kJ\,mol^{-1}$\n * $C-H$ single bond: $393\,kJ\,mol^{-1}$\n * $C ≡ O$ triple bond (in $CO$): $1074\,kJ\,mol^{-1}$\n\n* Resonance Structures\n * Definition: Structures that occur when it is possible to write two or more valid electron dot formulas that have the same number of electron pairs for a molecule or ion.\n * Example: Ozone ($O_3$). The actual bond length is an average (hybrid) of two resonance structures, not a rapid flipping between them.\n\n* Exceptions to the Octet Rule\n * Odd Number of Electrons: Molecules like Nitrogen Dioxide ($NO_2$) have a total of $17$ valence electrons; an octet for all atoms is impossible.\n * Electron Deficiency: Boron Trifluoride ($BF_3$) has only $6$ electrons around boron. It compensates by reacting with ammonia ($NH_3$) to form a coordinate covalent bond.\n * Expanded Octets: Phosphorus and Sulfur can expand to $10$ or $12$ electrons.\n * Phosphorus Pentachloride ($PCl_5$): $10$ valence electrons around P.\n * Sulfur Hexafluoride ($SF_6$): $12$ valence electrons around S.\n\n* Paramagnetism and Diamagnetism\n * Diamagnetic: Substances where all electrons are paired; weakly repelled by magnetic fields.\n * Paramagnetic: Substances with one or more unpaired electrons; strongly attracted to magnetic fields. Oxygen ($O_2$) and $NO_2$ are paramagnetic.\n * Ferromagnetism: Strong, permanent magnetism in iron, cobalt, and nickel due to alignment of ions ($Fe^{2+}$, $Co^{2+}$, $Ni^{2+}$).\n\n# 16.2 Bonding Theories\n\n* Molecular Orbital Theory\n * Assumes atomic orbitals overlap to produce molecular orbitals that apply to the whole molecule.\n * Bonding Orbital: A molecular orbital with energy lower than the atomic orbitals from which it formed. It is filled by electrons seeking lowest energy.\n * Antibonding Orbital: A molecular orbital with energy higher than the atomic orbitals; electrons here favor repulsion.\n\n* Bond Types (Sigma and Pi)\n * Sigma Bond ($σ$): Formed when two atomic orbitals combine to form a molecular orbital that is symmetrical along the axis connecting two nuclei. Formed by end-to-end overlap of $s$ or $p$ orbitals.\n * Pi Bond ($π$): Formed by the side-by-side overlap of atomic $p$ orbitals. Bonding electrons are found in sausage-shaped regions above and below the bond axis. Pi bonds are weaker than sigma bonds.\n\n* VSEPR Theory (Valence-Shell Electron-Pair Repulsion)\n * States that because electron pairs repel, molecular shape adjusts so valence-electron pairs are as far apart as possible.\n * Methane ($CH_4$): Tetrahedral shape, all angles are $109.5^∘$ (Tetrahedral Angle).\n * Ammonia ($NH_3$): Pyramidal shape. The unshared pair repels the bonding pairs more strongly, compressing angles to $107^∘$.\n * Water ($H_2O$): Bent shape. Two unshared pairs compress the angle to $105^∘$.\n * Carbon Dioxide ($CO_2$): Linear shape. No unshared pairs on carbon, angle is $180^∘$.\n\n* Hybridization\n * Several atomic orbitals mix to form the same number of equivalent hybrid orbitals.\n * $sp^3$ Hybridization: One $2s$ and three $2p$ orbitals mix. Results in $4$ orbitals at $109.5^∘$ (e.g., Methane).\n * $sp^2$ Hybridization: One $2s$ and two $2p$ orbitals mix. Results in $3$ orbitals at $120^∘$; the remaining $p$ orbital forms a pi bond (e.g., Ethene, $C_2H_4$).\n * $sp$ Hybridization: One $2s$ and one $2p$ orbital mix. Results in $2$ orbitals at $180^∘$ (linear); remaining two $p$ orbitals form two pi bonds (e.g., Ethyne/Acetylene, $C_2H_2$).\n\n# 16.3 Polar Bonds and Molecules\n\n* Bond Polarity\n * Nonpolar Covalent Bond: Electrons shared equally (e.g., $H_2, O_2$).\n * Polar Covalent Bond (Polar Bond): Electrons shared unequally. The more electronegative atom attracts electrons more strongly, acquiring a slightly negative charge ($δ^-$); the other acquires a slightly positive charge ($δ^+$).\n * Electronegativity Differences and Bond Types (Table 16.4):\n * $0.0-0.4$: Nonpolar covalent.\n * $0.4-1.0$: Moderately polar covalent.\n * $1.0-2.0$: Very polar covalent.\n * > 2.0: Ionic.\n\n* Polar Molecules\n * A molecule with two poles is a dipolar molecule or dipole.\n * In an electric field, polar molecules align with negative ends toward the positive plate and vice versa.\n * Bond polarity can cancel out depending on shape. Carbon Dioxide ($CO_2$) has polar bonds but is a nonpolar molecule because it is linear (polarities cancel). Water ($H_2O$) is polar because it is bent.\n\n* Intermolecular Attractions (Weakest to Strongest)\n 1. Van der Waals Forces: Collective name for weakest attractions.\n * Dispersion Forces: Weakest, caused by electron motion. Strength increases with number of electrons (Fluoroine gas vs. Iodine solid).\n * Dipole Interactions: Occur when polar molecules are attracted to one another (electrostatic attraction between opposite poles).\n 2. Hydrogen Bonds: Attractive forces where hydrogen covalently bonded to a very electronegative atom ($O, N, F$) is also weakly bonded to an unshared pair of another electronegative atom. Strength is about $5\%$ of an average covalent bond.\n\n* Intermolecular Attractions and Properties\n * Most molecular compounds have low melting points (usually below $300^∘\text{C}$) compared to ionic compounds.\n * Network Solids (Network Crystals): Solids where all atoms are covalently bonded to each other. They have very high melting points or decompose without melting.\n * Diamond: Network form of carbon; vaporizes at $3500^∘\text{C}$ and above.\n * Silicon Carbide ($SiC$): Melting point $2700^∘\text{C}$; used in grindstones.\n\n# Features and Labs\n\n* Society: Blocks and Bonds\n * Sunlight's ultraviolet (UV) radiation can break covalent bonds in skin cell molecules like DNA.\n * Damage to DNA can lead to uncontrolled cell division (cancer).\n * Melanin: Dark pigment that absorbs UV. Sun Protection Factor (SPF) measures effectiveness (e.g., SPF $6$ means $6$ hours of use equals $1$ hour of unprotected exposure).\n\n* Discover It! Shapes of Molecules (Lab)\n * Inflating clusters of $2, 3,$ and $4$ balloons modeling Linear, Trigonal Planar, and Tetrahedral shapes as they naturally push away from each other.\n\n* Small-Scale Lab: Paper Chromatography of Food Dyes\n * Separates dyes (Red No. $40$, Yellow No. $5$, Blue No. $1$) based on polarity using a $0.1\%\,NaCl$ solvent. Most polar dyes move fastest/highest.\n\n* Math/Geography/Computer Science Links\n * Topographic Maps: Scientists use electron density plots (similar to contour line maps) to pinpoint atom locations.\n * Molecular Modeling: Computer programs like RasMol allow chemists to view and manipulate 3D structures of complex proteins.\n * Oncologist: Medical career focused on using treatments to stop the spread of cancer cells.\n\n* Questions & Discussion\n * What creates magnetic fields in electrons? Electrons are spinning electric charges; opposite spins in pairs cancel out magnetism.\n * Why is methane unreactive? High bond dissociation energy for $C-C$ and $C-H$ bonds.\n * Why does a snowflake have distinctive geometry? Polar bonds in water molecules combined with air temperature/vapor determine the crystal shape (over $100$ crystals per snowflake).",

  "title": "Covalent Bonding: Exhaustive Study Notes" }