3: Ionic Compounds
- Page ID
- 86201
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)When you think of bonding, you may not think of ions or molecules. Like most of us, you probably think of bonding between people. Like people, molecules bond — and some bonds are stronger than others. It's hard to break up a mother and baby, or a molecule made up of one oxygen and two hydrogen atoms! A chemical bond is a force of attraction between atoms or ions. Bonds form when atoms share or transfer valence electrons. Valence electrons are the electrons in the outer energy level of an atom that may be involved in chemical interactions. Valence electrons are the basis of all chemical bonds.
- 3.1: Ions
- This page explains ions as charged particles from atom electron gain or loss. Cations are positive ions from lost electrons, while anions are negative ions from gained electrons. Naming conventions include adding "ion" for cations and modifying the element name to end in "ide" for anions. These concepts are essential for understanding chemical reactions and bonding.
- 3.2: Ions and the Octet Rule
- This page explains ion formation through electron transfer, emphasizing the octet rule for stable electron configurations. It details how main group metals lose electrons to become cations, such as sodium (Na+), while nonmetals gain electrons to form anions, such as chlorine (Cl–). The page also covers determining electron configurations and the periodic table's group behavior related to ion formation, along with potential octet rule exceptions.
- 3.3: Ions of Some Common Elements
- This page explains how to determine the charges of common ions using the periodic table, highlighting that elements in the same group form ions with similar charges due to their valence electrons. It notes that alkali metals form +1 ions, alkaline earth metals +2 ions, and halogens -1 ions, while transition metals, due to variable electron configurations, exhibit more complex charge patterns. Examples are provided to clarify these concepts.
- 3.4: Periodic Properties and Ion Formation
- This page explains ionization energy (IE) and electron affinity (EA), crucial for predicting ion formation trends. It notes that IE increases across a period and decreases down a group due to atomic size and nuclear charge, while exceptions can occur based on electron configurations. EA, reflecting the energy change when adding an electron, is more negative left to right across a period and less negative down a group.
- 3.5: Naming Monoatomic Ions
- This page discusses naming conventions for monoatomic ions in chemistry, highlighting the distinction between cations and anions. Cations adopt the element's name plus "ion," while anions end in "-ide." The stock system employs Roman numerals for variable charge elements, and the common system uses the suffixes "-ic" and "-ous." Examples are provided for ions like iron and copper, and the significance of common ions such as Na+ and Cl− for human health is emphasized.
- 3.6: Polyatomic Ions
- This page covers polyatomic ions, groups of covalently bonded atoms with an overall charge, emphasizing their structures, names, and formulas. It provides a table of common ions, noting that most are negatively charged, with two exceptions (hydronium and ammonium). The page explains the use of suffixes "-ate" and "-ite" to distinguish ions based on oxygen counts, stressing the importance of memorization for their names and formulas.
- 3.7: Ionic Bonds
- This page explains ionic bonds and compounds through electrostatic forces, detailing how oppositely charged ions attract. Using sodium chloride (NaCl) as an example, it describes the electron transfer that forms Na⁺ and Cl⁻ ions. The page emphasizes the three-dimensional arrangement of ions in solids to maximize attractive forces and minimize repulsive ones, highlighting the importance of charge magnitude and ion size in bond strength.
- 3.8: Formulas of Ionic Compounds
- This page covers the formation and representation of ionic compounds, detailing the importance of electric neutrality where positive and negative charges balance. It describes how monoatomic and polyatomic ions combine, using examples like NaCl, MgO, and Ca3(PO4)2.
- 3.9: Naming Ionic Compounds
- This page explains the naming conventions for ionic compounds, detailing how cations are named before anions without indicating charge. It differentiates between Type I (simple cations) and Type II (cations with variable charges) and provides examples using Stock and common naming systems. The page also includes a list of common ionic compounds, promoting recognition and naming skills in everyday contexts.
- 3.10: Some Properties of Ionic Compounds
- This page covers the physical properties of ionic compounds, noting their high melting points from strong ionic bonds, along with their hardness and brittleness due to ion arrangement. It explains that solid ionic compounds do not conduct electricity, while dissolved or melted ones can, due to the formation of mobile ions. Examples of ionic crystals showcase their diverse colors and shapes.
- 3.11: H⁺ and OH⁻ Ions - An Introduction to Acids and Bases
- This page covers the definitions of acids and bases, identifying acids as substances that produce hydrogen ions (\(\ce{H^{+}}\)) and bases as those that produce hydroxide ions (\(\ce{OH^{−}}\)) in water. It mentions well-known examples of acids (e.g., citric and acetic acid) and bases (e.g., sodium bicarbonate and sodium hydroxide).
Thumbnail: The crystal structure of sodium chloride, NaCl, a typical ionic compound. The purple spheres represent sodium cations, \(\ce{Na^{+}}\), and the green spheres represent chloride anions, \(\ce{Cl^{−}}\). (Public Domain; Benjah-bmm27 via Wikipedia)