4: Molecular Compounds
- Page ID
- 86208
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 4.1: Covalent Bonds
- This page covers the formation of covalent bonds between nonmetal atoms, focusing on the octet rule and electron sharing. It illustrates the concept using hydrogen molecules (H2) and bond lengths, while introducing Lewis structures to visualize bonding and lone pair electrons, exemplified by fluorine. The page also distinguishes between diatomic molecules, such as H2 and F2, and individual atoms, enhancing the understanding of molecular structures.
- 4.2: Covalent Bonds and the Periodic Table
- This page covers covalent bonding basics, detailing how to predict bond formation based on periodic table positioning and the octet rule. It includes examples of diatomic and polyatomic molecules like water and ammonia, and explains exceptions to the octet rule, such as odd-electron and electron-deficient molecules. The importance of understanding these concepts for grasping molecular structures and bonding behavior is emphasized, with illustrative examples and exercises provided for clarity.
- 4.3: Multiple Covalent Bonds
- This page covers the role of multiple bonds in fulfilling the octet rule for atoms in molecular structures, explaining single, double, and triple bonds. It provides examples with nitrogen, oxygen, and carbon, and guides readers through drawing Lewis structures for molecules like hydrogen cyanide, ethane, and acetylene. The page details the calculation of valence electrons and the arrangement necessary for octet completion.
- 4.4: Coordinate Covalent Bonds
- This page discusses coordinate covalent bonds, where one atom donates both electrons in a shared pair, distinguishing them from standard covalent bonds. It uses carbon monoxide (CO) to illustrate this concept with its triple bond, which includes a coordinate bond. The page also addresses the reactivity of electron-deficient molecules like BF3, which can accept lone pairs from donors like NH3 to form coordinate covalent bonds.
- 4.5: Characteristics of Molecular Compounds
- This page explains the distinctions between ionic and molecular compounds, highlighting how their bonding types affect their physical properties. Ionic compounds, formed by electron transfer, exhibit higher melting and boiling points, good conductivity in solution, and high water solubility. In contrast, molecular compounds, created by covalent bonds, have lower melting and boiling points, poor conductivity, and variable solubility, with their properties also affected by molecular polarity.
- 4.6: Molecular Formulas and Lewis Structures
- This page covers methods of representing molecules in chemistry, focusing on the need for a "universal language" among chemists. It explains molecular formulas, structural formulas, and various models like ball-and-stick and space-filling models, which aid in visualizing molecular structures. Additionally, it mentions condensed structural formulas as a streamlined way of representing molecules.
- 4.7: Drawing Lewis Structures
- This page provides a seven-step guide for drawing Lewis structures of molecules and ions, detailing the process from atom arrangement to the formation of multiple bonds. It includes examples like H2O and CH2O to demonstrate these steps and highlights the importance of electronegativity in selecting central atoms. Additionally, the page offers practice exercises to reinforce the concepts learned.
- 4.8: The Shapes of Molecules
- This page covers the valence shell electron pair repulsion (VSEPR) theory, which predicts molecular structures by arranging electron groups around a central atom to minimize repulsion. It differentiates between electron group geometry and molecular structure, providing a method using Lewis diagrams. Examples illustrate VSEPR applications, including a focus on molecules with multiple central atoms, like methyl isocyanate, highlighting their various shapes.
- 4.9: Polar Covalent Bonds and Electronegativity
- This page covers electronegativity, bond polarity, and the distinction between ionic and covalent bonds, highlighting electron transfer and sharing. It introduces Pauling's electronegativity scale, emphasizes trends in the periodic table, and explains how electronegativity differences impact bond types. Notable figures like Linus Pauling are discussed for their contributions to chemistry and societal issues, despite later controversies regarding their claims, ensuring their legacy endures.
- 4.10: Polar Molecules
- This page discusses covalent bond polarity, outlining how single bonds influence molecular polarity through examples of polar (HF, water) and nonpolar (CO2, CH4) molecules. It highlights the significance of geometric arrangement in determining polarity, provides steps for identifying polar molecules, and covers the properties of polar molecules, including their behavior in electric fields and solubility in various solvents.
- 4.11: Naming Binary Molecular Compounds
- This page explains the naming conventions for binary molecular compounds formed from two nonmetals, detailing that the first element keeps its name while the second adopts the -ide suffix. Numerical prefixes indicate atom quantities, and a table provides this information along with noted exceptions. Examples are given for naming and deriving molecular formulas, highlighting common names like water and ammonia.
Thumbnail: Covalently bonded hydrogen and carbon in a w:molecule of methane. (CC BY-SA 2.5; DynaBlast via Wikipedia)