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9: Covalent Bonding

Last updated

Mar 21, 2025
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- 8.12: Alloys
- 9.1: Chemical Bond

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Molecules are groups of atoms that behave as a single unit. Some elements exist as molecules: hydrogen, oxygen, sulfur, and so forth. There are rules that can express a unique name for any given molecule, and a unique formula for any given name.

9.1: Chemical Bond
This page explains chemical bonds as attractive forces between atoms or ions formed by sharing or transferring valence electrons to achieve stability. It describes three main types: covalent bonds (electron sharing among nonmetals), ionic bonds (formed between oppositely charged ions creating crystals), and metallic bonds (in metals where positive ions are surrounded by free-flowing valence electrons).
9.2: Covalent Bond
This page explains covalent bonds as attractions between nonmetal atoms sharing valence electrons, which form molecules and covalent compounds like water. These bonds stabilize atoms by filling their outer energy levels. Diatomic molecules, such as oxygen, consist of two atoms sharing multiple electron pairs. Covalent bonds are typically stronger than ionic bonds, and covalent compounds do not conduct electricity.
9.3: Molecular Compounds
This page discusses covalent compounds formed by nonmetals sharing valence electrons, highlighting their properties such as easy combustion, poor water solubility, and lack of electrical conductivity. Naming conventions include placing the left-most element first, using the suffix -ide for the second, and prefixes for quantities. An example is dinitrogen trioxide (N2O3).
9.4: Energy and Covalent Bond Formation
This page discusses the bonding differences between BeCl2 and LiCl, highlighting that LiCl forms ionic bonds while BeCl2 involves covalent bonds. It also covers the nature of molecular compounds, which are nonmetals that achieve stability by bonding and lowering potential energy. The page explains how atoms such as hydrogen stabilize through shared electrons, achieving ideal bond distances at minimum potential energy.
9.5: Lewis Electron-Dot Structures
This page explains cholesterol's molecular structure (C27H46O) and its detailed atomic arrangement. It describes how Lewis electron-dot structures represent valence electrons and covalent bonds through shared pairs. The concept of the octet rule is introduced, emphasizing that atoms share electrons for stability, with hydrogen forming H2 to resemble helium. Other atoms also share electrons to achieve a complete set of eight valence electrons.
9.6: Single Covalent Bonds
This page explains molecular formation through atom bonding, focusing on covalent bonds where orbitals overlap to share electrons. It gives examples like hydrogen ( $\ce{H_2}$ ) and fluorine ( $\ce{F_2}$ ), discusses the octet rule's role in achieving stable electron configurations, and illustrates principles with diagrams. It emphasizes the importance of electron sharing and lone pairs in molecular structures, particularly in compounds like water ( $\ce{H2O}$ ).
9.7: Multiple Covalent Bonds
This page outlines the process of managing leftover electrons in Lewis structures for covalent compounds and the necessity of double and triple bonds to satisfy the octet rule. It uses ethene ( $\ce{C2H4}$ ) as an example of double bonding and nitrogen ( $\ce{N2}$ ) for triple bonding. The text concludes with a summary of these bonding types and includes review questions for further understanding.
9.8: Coordinate Covalent Bond
This page explains the analogy of sharing, linking personal experiences of sharing toys to the sharing of electrons in chemistry. It covers coordinate covalent bonds, particularly in carbon monoxide ( $\ce{CO}$ ), and describes how a triple covalent bond involves one coordinate bond. The text emphasizes that this bond is just as strong as traditional covalent bonds and concludes with a review section that questions bond origins, structural inaccuracies, and comparative strengths.
9.9: Covalent Bonding in Polyatomic Ions
This page explores the difficulties the U.S. Supreme Court encounters in law interpretation, likening these to the systematic method in chemistry for drawing Lewis structures of polyatomic ions. It explains how to create these structures by calculating total valence electrons, accounting for charges, and satisfying the octet rule, using the ammonium ion ( $\ce{NH_4^+}$ ) and sulfate ion ( $\ce{SO_4^{2-}}$ ) as examples.
9.10: Resonance
This page discusses resonance in chemistry, using ozone ( $O_3$ ) to illustrate how some molecules can have multiple valid Lewis structures. It explains that the actual molecular structure is an average of these resonance structures, resulting in bonds that can be described as "one and a half" bonds, such as the bond lengths in ozone. The concept applies to other polyatomic ions, like the nitrate ion ( $NO_3^-$ ), which also exhibit resonance.
9.11: Exceptions to the Octet Rule
This page explains exceptions to the octet rule in chemistry, categorizing them into incomplete octets, odd-electron molecules, and expanded octets.
9.12: Bond Energy
This page discusses smog formation, mainly attributed to nitrogen compounds like $\ce{NO}_x$ from high-temperature combustion in car engines. It explains nitrogen's inertness as $\ce{N_2}$ and its reactivity upon breaking triple bonds. Additionally, it highlights the importance of bond energy in determining the stability and reactivity of various compounds.
9.13: VSEPR Theory
This page discusses the water molecule's bent shape, which enhances its polarity and boiling point, crucial for supporting life. It highlights the historical challenge in explaining this shape compared to carbon dioxide. In 1956, VSEPR theory was introduced by R.J. Gillespie and R.S. Nyholm, providing a method to predict molecular geometry by considering electron pair repulsion, aiding in the understanding of molecular structure based on bonding and lone electron pairs.
9.14: Molecular Shapes- No Lone Pairs on Central Atoms
This page discusses the electroscope, a device used to study charge through the movement of electrons, leading to repelling leaves. It also covers molecular structures using VSEPR theory, showcasing examples such as Beryllium Hydride and Carbon Dioxide (both linear), Boron Trifluoride (trigonal planar), and Methane (tetrahedral with 109.5° bond angles).
9.15: Molecular Shapes - Lone Pair(s) on Central Atom
This page explains how lone pair electrons influence the molecular geometry of compounds, highlighting examples like ammonia (NH₃) and water (H₂O) with their trigonal pyramidal and bent shapes, respectively. It also discusses sulfur tetrafluoride (SF₄) and its distorted tetrahedral geometry.
9.16: Bond Polarity
This page explains bond polarity in chemistry, detailing how electronegativity affects electron sharing or transfer between atoms. It describes that large electronegativity differences result in ionic bonds, while smaller differences lead to covalent bonds. Nonpolar covalent bonds feature equal electron sharing, whereas polar covalent bonds exhibit unequal sharing, creating partial charges. The page highlights the relevance of bond polarity in applications like soap cleaning.
9.17: Polar Molecules
This page discusses the creation of ultracold polar molecules using lasers to excite $\ce{Rb}$ and $\ce{K}$ atoms into charged $\ce{Rb-K}$ compounds at near absolute zero. It highlights the significance of molecular geometry in determining polarity and its implications for aligning molecules in an electric field. The understanding of these properties is crucial for developing new reactions and materials.
9.18: Van der Waals Forces
This page discusses the use of liquid nitrogen in MRI to cool superconducting magnets and outlines the significance of intermolecular forces like van der Waals, dipole-dipole, and London dispersion forces in molecular interactions.
9.19: Hydrogen Bonding
This page discusses the differences in boiling points and molecular weights of ammonia and nitrogen, explaining why ammonia has a higher boiling point due to hydrogen bonding. It details how hydrogen bonds, occurring between hydrogen and highly electronegative atoms (N, O, F), contribute to water's unique properties, including its liquid state at room temperature and ice's lower density. The text highlights the importance of hydrogen bonding in biological structures.
9.20: Physical Properties and Intermolecular Forces
This page discusses the two major forms of carbon: diamond, known for its hardness, and graphite, which is softer. It emphasizes that a compound's properties are influenced by chemical bonding, affecting melting and boiling points. Molecular compounds typically have lower melting and boiling points and poor electrical conductivity, with solubility in water influenced by hydrogen bonding. Covalent network solids like diamond have strong bonds, necessitating extreme temperatures for vaporization.
9.21: Valence Bond Theory
This page covers valence bond theory, detailing how covalent bonds arise from the overlap of atomic orbitals as atoms approach each other, leading to stable bonds at specific distances. It uses examples like $\ce{H_2}$ and $\ce{F_2}$ to illustrate key concepts. Review questions focus on electron positions in orbitals, covalent bond formation, and the necessity of overlapping orbitals being of the same type.
9.22: Hybrid Orbitals - sp³
This page explores hybridization in chemistry, centering on carbon's bonding in methane ( $\ce{CH_4}$ . Despite its electron configuration indicating it should form only two bonds, carbon actually forms four through $sp^3$ hybridization, which mixes one $s$ and three $p$ orbitals. This process results in four equivalent hybrid orbitals, accounting for methane's tetrahedral geometry and bond angles.
9.23: Hybrid Orbitals - sp and sp²
This page explains hybridization in chemistry, drawing a metaphor between paired electrons and the lovers Romeo and Juliet, who bond only when unpaired. It details $sp$ hybridization with beryllium hydride $\left( \ce{BeH_2} \right)$ for linear bonding and $sp^2$ hybridization with boron trifluoride $\left( \ce{BF_3} \right)$ for a trigonal planar structure. The page underscores the role of hybridization in allowing molecules to achieve specific geometries for optimal bonding.
9.24: Sigma and Pi Bonds
This page explains the hybridization of carbon atoms in molecules with double and triple bonds, using ethene ( $\ce{C2H4}$ ) and ethyne ( $\ce{C2H2}$ ) as examples. Ethene has $sp^2$ hybridization, yielding a planar structure with one sigma and one pi bond. In contrast, ethyne has $sp$ hybridization, resulting in a linear shape with one sigma and two pi bonds.

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