5: Chemical Bonding and Nomenclature
- Page ID
- 568630
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 5.1: A Molecular View of Elements and Compounds
- This page outlines the classification of substances into four categories: atomic elements (individual atoms), molecular elements (two or more bonded atoms), ionic compounds (metal and non-metal bond), and covalent (molecular) compounds (non-metals only). It provides definitions, examples, and exercises to enhance understanding of these classifications.
- 5.2: 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.
- 5.3: 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.
- 5.4: Lewis Structures of Covalent Compounds- 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.
- 5.5: Rules for Drawing Lewis Structures
- 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.
- 5.6: 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.
- 5.7: 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.
- 5.8: Chemical Nomenclature
- This page discusses the importance of nomenclature in chemistry, emphasizing the need for systematic naming of compounds to prevent ambiguity and enhance communication. It highlights the goal of clarity and uniqueness in names, alongside the reflection of a compound's structure or chemistry. While common names may be acceptable in certain contexts, the ideal chemical name should communicate the chemical formula.
- 5.9: Writing Formulas for Ionic Compounds
- This page covers the formulation and identification of ionic compounds, detailing their crystal lattice structure and the derivation of empirical formulas through charge balancing of cations and anions. It introduces the crisscross method for formula writing and the role of polyatomic ions. Specific examples like potassium sulfate and magnesium carbonate illustrate the concepts.
- 5.10: Naming Ionic Compounds
- This page covers the naming rules for ionic compounds, detailing the nomenclature of monatomic cations and anions. Cations are named with "ion" added, while anions use the element stem plus "-ide." The Stock system employs Roman numerals for cations with variable charges, and traditional Latin names are also utilized. It emphasizes balancing positive and negative charges, lists examples, and introduces polyatomic ions.
- 5.11: Naming Covalent Compounds
- This page discusses naming covalent compounds, which are formed by nonmetal atoms sharing electrons, contrasting them with ionic compounds based on electrostatic attraction. It explains the naming convention for binary covalent (molecular) compounds using Greek prefixes and the -ide suffix, citing examples like water and ammonia.
- 5.12: Naming Acids
- This page details how acids release hydrogen ions in water and categorizes them into binary acids and oxyacids, providing naming conventions based on anion types along with examples. The page also addresses how to write formulas for acids, focusing on maintaining charge neutrality.
- 5.13: Nomenclature Summary
- This page offers a brief overview of chemical nomenclature in the form of a visual flowchart.