10: Chemical Bonding
- Page ID
- 47424
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)How do atoms make compounds? Typically they join together in such a way that they lose their identities as elements and adopt a new identity as a compound. These joins are called chemical bonds. But how do atoms join together? Ultimately, it all comes down to electrons. Before we discuss how electrons interact, we need to introduce a tool to simply illustrate electrons in an atom.
- 10.2: Representing Valence Electrons with Dots
- This page explains Lewis electron dot diagrams, which depict valence electrons as dots around an element's symbol, with conventions for dot arrangement based on electron configurations. It illustrates examples such as hydrogen with one dot and helium with two, covering elements up to neon. Ion diagrams are also discussed, showing adjustments in dot numbers for cations and anions.
- 10.3: Lewis Structures of Ionic Compounds- Electrons Transferred
- This page covers the octet rule, stating that atoms are most stable with eight electrons in their valence shell. It explains how ions form via electron loss (cations) or gain (anions) to achieve stability, and introduces ionic bonding through electron transfer, using sodium chloride and magnesium oxide as examples. The balance of charges in ionic compounds is emphasized, along with the importance of charge magnitudes and ion sizes in affecting ionic bond strength.
- 10.4: Covalent Lewis Structures- Electrons Shared
- This page explains covalent bonding, highlighting how atoms share electrons for valence shell completion, unlike ionic bonding. It uses hydrogen and fluorine examples with Lewis dot diagrams to illustrate bond formation, covering single, double, and triple bonds, especially in molecules like ethene and nitrogen. The goal is to help readers define covalent bonds and depict their formation accurately.
- 10.5: Writing Lewis Structures for Covalent Compounds
- This page covers the creation of Lewis structures for covalent compounds and polyatomic ions, detailing steps such as counting valence electrons, arranging atoms, and ensuring octets. It provides examples like H2O and CH2O, along with polyatomic ions, discussing how charges influence valence counting.
- 10.6: Resonance - Equivalent Lewis Structures for the Same Molecule
- This page discusses resonance in molecules using ozone (\(\ce{O3}\) ) as an example. It explains that resonance arises from multiple valid Lewis structures, leading to uniform bond characteristics rather than distinct single or double bonds. In ozone, each oxygen bond is considered a "one and a half" bond, reflecting an average of resonance structures. The concept extends to polyatomic ions, such as the nitrate ion (\(\ce{NO3^-}\)), where bonds are also averaged.
- 10.7: Predicting the Shapes of Molecules
- This page explains the shapes of simple molecules using valence shell electron pair repulsion (VSEPR) theory, detailing how electron repulsion influences molecular geometry. It categorizes shapes into linear, trigonal planar, and tetrahedral formations, depending on the arrangement of electron groups around a central atom.
- 10.8: Electronegativity and Polarity - Why Oil and Water Do not Mix
- This page explores the distinction between polar and nonpolar compounds, focusing on electronegativity's impact on bond polarity and the classification of bonds. It explains that nonpolar covalent bonds feature equal electron sharing, while polar bonds exhibit unequal sharing, leading to partial charges. The discussion extends to molecular polarity, emphasizing how geometry affects it, and outlines steps for determining polarity through Lewis structures and VSEPR theory.