Chemistry for Changing Times - Chapter 4: Chemical Bonds

Chemical Bonds

Introduction to Chemical Bonds

Carbon exists commonly as soot, which can transform into diamonds under high temperature and pressure. This transformation is a direct result of changes in the chemical bonds holding the atoms together.

Stable Electron Configurations

Fact: Noble gases (e.g., helium, neon, argon) are chemically inert, meaning they undergo few chemical reactions.
Theory: This inertness is attributed to their electron structures. All noble gases, except helium, possess an octet of electrons (eight valence electrons) in their outermost electron shell.
Deduction: Elements become less reactive when they achieve an electron structure similar to that of a noble gas.
- Example (Sodium): Sodium (Na) can lose one valence electron. After losing this electron, its remaining core electrons are configured like the noble gas neon ( $Ne$ ), resulting in a $Na^+$ ion.
- Example (Chlorine): Chlorine (Cl) can gain one electron. By doing so, its electron structure becomes like the noble gas argon ( $Ar$ ), forming a $Cl^-$ ion with $18$ electrons ( $17$ protons $+ 1 $electron gained).</li></ul></li></ul><h4 id="83e69278-e34d-428f-a509-8fef44c0e998" data-toc-id="83e69278-e34d-428f-a509-8fef44c0e998" collapsed="false" seolevelmigrated="true">Lewis (Electron Dot) Symbols</h4><ul><li>Developed by G. N. Lewis, this method visually represents valence electrons as dots around the atomic symbol.</li><li>Table 4.1 Lewis Dot Symbols for Selected Main Group Elements:<ul><li>Group 1A:$ Hullet $,$ Liullet $,$ Naullet $,$ Kullet $,$ Rbullet $,$ Csullet $</li><li>Group 2A:$ \bullet Beullet $,$ \bullet Mgullet $,$ \bullet Caullet $,$ \bullet Srullet $,$ \bullet Baullet $</li><li>Group 3A:$ \cdot B\cdot $,$ \cdot Al\cdot $</li><li>Group 4A:$ \cdot C\cdot $,$ \cdot Si\cdot $</li><li>Group 5A:$ :.N\cdot $,$ :.P\cdot $</li><li>Group 6A:$ :\ddot O\cdot $,$ :\ddot S\cdot $</li><li>Group 7A:$ :\ddot{Cl}\cdot $,$ :\ddot{Br}\cdot $</li><li>Noble Gases:$ He: $,$ :\ddot{Ne}: $</li><li>Note: While most elements strive for an octet, helium only needs two electrons for a stable configuration.</li></ul></li></ul><h4 id="3de8e28c-05a8-40b9-a844-3f6773977fa8" data-toc-id="3de8e28c-05a8-40b9-a844-3f6773977fa8" collapsed="false" seolevelmigrated="true">Ionic Bonds</h4><ul><li>Formation: When sodium reacts with chlorine, the sodium atom transfers an electron to the chlorine atom. This forms a$ Na^+ $ion (like neon) and a$ Cl^- $ion (like argon).</li><li>Nature of Attraction:$ Na^+ $ions and$ Cl^- $ions, having opposite charges, attract each other. This electrostatic attraction constitutes an ionic bond.</li><li>Structure: Ionic compounds are held together by these opposite electrostatic charges and exist as a crystal lattice structure.</li><li>Atoms and Ions: Distinct Differences: An atom is electrically neutral, while an ion carries an electrical charge due to the gain or loss of electrons. For example, calcium metal (Ca) is distinct from calcium ions ($ Ca^{2+} $) in dietary supplements, and iron metal (Fe) is distinct from iron ions ($ Fe^{2+} $or$ Fe^{3+} $) in supplements.</li></ul><h4 id="40dcbbbf-7063-48fe-b919-713c6aebe239" data-toc-id="40dcbbbf-7063-48fe-b919-713c6aebe239" collapsed="false" seolevelmigrated="true">The Octet Rule</h4><ul><li>In chemical reactions, atoms generally tend to gain, lose, or share electrons to achieve eight valence electrons. This fundamental principle is known as the octet rule. It is particularly valid for elements in the first and second primary shells.</li><li>Metals: Metals typically lose electrons to attain the electron structure of the previous noble gas, forming positive ions (cations).</li><li>Nonmetals: Nonmetals generally gain electrons to achieve the electron structure of the next noble gas, forming negative ions (anions).</li></ul><h4 id="d05b55aa-54d3-48be-8f92-678c3d7dcd02" data-toc-id="d05b55aa-54d3-48be-8f92-678c3d7dcd02" collapsed="false" seolevelmigrated="true">Symbols and Names for Some Simple (Monatomic) Ions</h4><ul><li>A comprehensive list of common monatomic ions and their names is provided in Table 4.2.<ul><li>Group 1A (Cations): Hydrogen ($ H^+ $, Hydrogen ion), Lithium ($ Li^+ $, Lithium ion), Sodium ($ Na^+ $, Sodium ion), Potassium ($ K^+ $, Potassium ion).</li><li>Group 2A (Cations): Magnesium ($ Mg^{2+} $, Magnesium ion), Calcium ($ Ca^{2+} $, Calcium ion).</li><li>Group 3A (Cations): Aluminum ($ Al^{3+} $, Aluminum ion).</li><li>Group 5A (Anions): Nitrogen ($ N^{3-} $, Nitride ion), Phosphorus ($ P^{3-} $, Phosphide ion).</li><li>Group 6A (Anions): Oxygen ($ O^{2-} $, Oxide ion), Sulfur ($ S^{2-} $, Sulfide ion).</li><li>Group 7A (Anions): Fluorine ($ F^- $, Fluoride ion), Chlorine ($ Cl^- $, Chloride ion), Bromine ($ Br^- $, Bromide ion), Iodine ($ I^- $, Iodide ion).</li><li>Transition Metals (Variable Charge Cations):<ul><li>Copper ($ Cu^+ $, Copper(I) ion or cuprous ion;$ Cu^{2+} $, Copper(II) ion or cupric ion).</li><li>Silver ($ Ag^+ $, Silver ion).</li><li>Zinc ($ Zn^{2+} $, Zinc ion).</li><li>Iron ($ Fe^{2+} $, Iron(II) ion or ferrous ion;$ Fe^{3+} $, Iron(III) ion or ferric ion).</li></ul></li></ul></li></ul><h4 id="10880472-5941-47fc-972a-42f2612087ab" data-toc-id="10880472-5941-47fc-972a-42f2612087ab" collapsed="false" seolevelmigrated="true">Formulas and Names of Binary Ionic Compounds</h4><ul><li>Cation Charge: For representative elements, the charge of a cation is equal to its group (family) number. The cation's name is simply the element's name.<ul><li>Examples:$ Na^+ $= sodium ion,$ Mg^{2+} $= magnesium ion.</li></ul></li><li>Anion Charge: For representative elements, the charge of an anion is equal to the group number minus eight ($ Group # - 8 $). The anion's name is the element's root name plus the suffix "-ide".<ul><li>Examples:$ Cl^- $= chloride ion,$ O^{2-} $= oxide ion.</li></ul></li><li>Naming Binary Ionic Compounds: Simply name the ions in order (cation first, then anion).<ul><li>Examples:$ NaCl $= sodium chloride,$ MgO $= magnesium oxide.</li></ul></li><li>Transition Metals: Many transition metals can exhibit more than one ionic charge. Roman numerals enclosed in parentheses are used to denote the charge of these ions in their names.<ul><li>Examples:$ Fe^{2+} $= iron(II) ion,$ Fe^{3+} $= iron(III) ion,$ Cu^{2+} $= copper(II) ion,$ Cu^+ $= copper(I) ion.</li></ul></li></ul><h4 id="45669055-873a-4726-8fb7-d3e8b8259319" data-toc-id="45669055-873a-4726-8fb7-d3e8b8259319" collapsed="false" seolevelmigrated="true">Covalent Bonds</h4><ul><li>Formation: Many nonmetallic elements react by sharing electrons rather than gaining or losing them.</li><li>Definition: When two atoms share a pair of electrons, a covalent bond is formed.</li><li>Types: Atoms can share one, two, or three pairs of electrons, forming single, double, and triple bonds, respectively.</li></ul><h4 id="d5e7da9d-b5f8-4617-8621-eb536dbd3776" data-toc-id="d5e7da9d-b5f8-4617-8621-eb536dbd3776" collapsed="false" seolevelmigrated="true">Names of Binary Covalent Compounds</h4><ul><li>Binary covalent compounds are named using prefixes to indicate the number of atoms of each element.</li><li>Naming Convention: A binary covalent compound has two parts to its name:<ol><li>First name: <code>prefix + name of first element</code>. If the first element has only one atom, the prefix "mono-" is dropped.</li><li>Second name: <code>prefix + root name of second element + suffix -ide</code>.</li></ol></li><li>Examples:<ul><li>$ SBr_4 $= Sulfur tetrabromide</li><li>$ P2O3 $= Diphosphorus trioxide</li></ul></li></ul><h4 id="b0454f83-f142-4605-8da2-9396bc625a0d" data-toc-id="b0454f83-f142-4605-8da2-9396bc625a0d" collapsed="false" seolevelmigrated="true">Electronegativity and Polar Covalent Bonds</h4><ul><li>Electronegativity: This is a measure of an atom's attraction for the bonding electrons in a chemical bond.<ul><li>Periodic Trend Example (Table 4.3): From left to right across a period, electronegativity generally increases (e.g., Li ($ 1.0 $) to F ($ 4.0 $)). The most electronegative element is Fluorine ($ 4.0 $).</li></ul></li><li>Polar Covalent Bond: When two atoms with differing electronegativities form a bond, the bonding electrons are pulled closer to the atom with the higher electronegativity. This unequal sharing of electrons creates a separation of charge within the bond, resulting in a polar covalent bond.</li><li>Representing Bond Polarity:<ul><li>Using partial charges:$ $\delta^+ $(delta plus) on the less electronegative atom and$ $\delta^- $(delta minus) on the more electronegative atom.</li><li>Using an arrow: An arrow pointing towards the more electronegative atom, with a cross at the tail representing the positive end (e.g.,$ A \text{---}> B $where B is more electronegative).</li></ul></li><li>Determining Bond Type by Electronegativity Difference: A general rule of thumb can classify bond types:<ul><li>If$ | ext{EN}1 - ext{EN}2| \approx 0 $: Nonpolar Covalent</li><li>If$ 0 < | ext{EN}1 - ext{EN}2| \le 0.4 $: Nonpolar Covalent</li><li>If$ 0.4 < | ext{EN}1 - ext{EN}2| \le 1.7 $: Polar Covalent</li><li>If$ | ext{EN}1 - ext{EN}2| > 1.7 $: Ionic</li></ul></li></ul><h4 id="3d116c2a-252d-4744-a0f4-3657d1ac1433" data-toc-id="3d116c2a-252d-4744-a0f4-3657d1ac1433" collapsed="false" seolevelmigrated="true">Polyatomic Ions</h4><ul><li>Definition: Polyatomic ions are groups of covalently bonded atoms that collectively carry an electrical charge.</li><li>Table 4.4 Some Common Polyatomic Ions:<ul><li>Charge 1+: Ammonium ion ($ NH4^+ $), Hydronium ion ($ H3O^+ $).</li><li>Charge 1-: Hydrogen carbonate (bicarbonate) ion ($ HCO3^- $), Hydrogen sulfate (bisulfate) ion ($ HSO4^- $), Acetate ion ($ CH3CO2^- $or$ C2H3O2^- $), Nitrite ion ($ NO2^- $), Nitrate ion ($ NO3^- $), Cyanide ion ($ CN^- $), Hydroxide ion ($ OH^- $), Dihydrogen phosphate ion ($ H2PO4^- $), Permanganate ion ($ MnO4^- $).</li><li>Charge 2-: Carbonate ion ($ CO3^{2-} $), Sulfate ion ($ SO4^{2-} $), Chromate ion ($ CrO4^{2-} $), Hydrogen (monohydrogen) phosphate ion ($ HPO4^{2-} $), Oxalate ion ($ C2O4^{2-} $), Dichromate ion ($ Cr2O7^{2-} $).</li><li>Charge 3-: Phosphate ion ($ PO_4^{3-} $).</li></ul></li><li>Writing Formulas with Polyatomic Ions: When a compound contains more than one polyatomic ion of a particular type, parentheses are used around the polyatomic ion's formula to indicate the proper number.<ul><li>Example: For calcium nitrate, with$ Ca^{2+} $and$ NO3^- $, the formula is$ Ca(NO3)_2 $.</li></ul></li><li>Naming Compounds with Polyatomic Ions: Simply name the ions in order (cation first, then anion).<ul><li>Example:$ (NH4)2SO_4 $= Ammonium sulfate.</li></ul></li></ul><h4 id="e90d8aba-d470-4da8-b885-fd38acfffb61" data-toc-id="e90d8aba-d470-4da8-b885-fd38acfffb61" collapsed="false" seolevelmigrated="true">Rules for Sketching Lewis Structures</h4><ul><li>Step 1: Count the grand total of all valence electrons from all atoms in the molecule or ion.</li><li>Step 2: Sketch a skeletal structure, connecting atoms with single bonds. The least electronegative atom is typically the central atom (except for hydrogen).</li><li>Step 3: Place remaining electrons as lone pairs around the outer (terminal) atoms first, ensuring each outer atom (except hydrogen) satisfies the octet rule (or duo-rule for hydrogen).</li><li>Step 4: Subtract the electrons used so far from the total number of valence electrons. Place any remaining electrons around the central atom as lone pairs.</li><li>Step 5: If the central atom lacks an octet (and has less than 8 electrons), move one or more lone pairs from an outer atom to form double or triple bonds with the central atom until the central atom achieves an octet.</li><li>HONC Rules (Implicit in Table 4.5 - Number of Bonds Formed by Selected Elements):<ul><li>Hydrogen ($ H $) forms$ 1 $bond (e.g.,$ H-H $,$ H-Cl $).</li><li>Oxygen ($ O $) typically forms$ 2 $bonds and$ 2 $lone pairs (e.g.,$ H-O-H $).</li><li>Nitrogen ($ N $) typically forms$ 3 $bonds and$ 1 $lone pair (e.g.,$ H-N-H $).</li><li>Carbon ($ C $) typically forms$ 4 $bonds and no lone pairs (e.g.,$ H_3C-H $).</li></ul></li></ul><h4 id="0fde8c6e-a9cf-492f-ab55-24bd089a12cb" data-toc-id="0fde8c6e-a9cf-492f-ab55-24bd089a12cb" collapsed="false" seolevelmigrated="true">Odd Electron Molecules: Free Radicals</h4><ul><li>Definition: An atom or molecule that possesses one or more unpaired electrons is called a free radical.</li><li>Examples:$ NO $,$ NO2 $,$ ClO2 $.</li></ul><h4 id="938ec827-9291-4ee2-9747-2ee6b76da1c8" data-toc-id="938ec827-9291-4ee2-9747-2ee6b76da1c8" collapsed="false" seolevelmigrated="true">Molecular Shapes: The VSEPR Theory</h4><ul><li>VSEPR Theory (Valence Shell Electron Pair Repulsion): This theory predicts the three-dimensional shape of molecules and polyatomic ions based on the principle that electron pairs (both bonding and lone pairs) around a central atom will orient themselves as far apart as possible to minimize repulsion.</li><li>Electron Set Arrangements and Optimal Angles:<ul><li>Two electron sets: Linear geometry, optimal angle$ 180 extdegree $.</li><li>Three electron sets: Trigonal planar geometry, optimal angle$ 120 extdegree $.</li><li>Four electron sets: Tetrahedral geometry, optimal angle$ 109.5 extdegree $(in 3D).</li></ul></li><li>Table 4.6 Bonding and the Shapes of Molecules (Summary):<ul><li>2 Electron Sets:<ul><li>$ 2 $Bonded Atoms,$ 0 $Lone Pairs: Linear. Examples:$ BeCl2 $,$ HgCl2 $,$ CO_2 $,$ HCN $.</li></ul></li><li>3 Electron Sets:<ul><li>$ 3 $Bonded Atoms,$ 0 $Lone Pairs: Trigonal Planar. Examples:$ BF3 $,$ AlBr3 $,$ CH_2O $.</li><li>$ 2 $Bonded Atoms,$ 1 $Lone Pair: Bent. Examples:$ SO2 $,$ O3 $.</li></ul></li><li>4 Electron Sets:<ul><li>$ 4 $Bonded Atoms,$ 0 $Lone Pairs: Tetrahedral. Examples:$ CH4 $,$ CBr4 $,$ SiCl_4 $.</li><li>$ 3 $Bonded Atoms,$ 1 $Lone Pair: Trigonal Pyramidal. Examples:$ NH3 $,$ PCl3 $.</li><li>$ 2 $Bonded Atoms,$ 2 $Lone Pairs: Bent. Examples:$ H2O $,$ H2S $,$ SCl_2 $.</li></ul></li></ul></li></ul><h4 id="a32bc087-6434-46bc-b454-3941943e44a0" data-toc-id="a32bc087-6434-46bc-b454-3941943e44a0" collapsed="false" seolevelmigrated="true">Shapes and Properties: Polar and Nonpolar Molecules</h4><ul><li>For a molecule to be classified as polar, two essential conditions must be met:<ol><li>The molecule must contain polar bonds (due to differences in electronegativity between bonded atoms).</li><li>The polar bonds must be arranged in such a way that there is an overall net separation of charge across the molecule. If the bond polarities cancel each other out due to molecular symmetry, the molecule will be nonpolar.</li></ol></li><li>Examples illustrate this concept (as depicted in figures):<ul><li>Methane ($ CH4 $): Although$ C-H $bonds have slight polarity, the tetrahedral arrangement (symmetric) causes the bond polarities to cancel, making$ CH4 $a nonpolar molecule.</li><li>Ammonia ($ NH3 $):$ N-H $bonds are polar, and the trigonal pyramidal shape (asymmetric, with a lone pair on nitrogen) results in an overall net dipole moment, making$ NH3 $a polar molecule.</li><li>Water ($ H2O $):$ O-H $bonds are highly polar, and the bent shape (asymmetric, with two lone pairs on oxygen) results in a significant net dipole moment, making$ H2O $a highly polar molecule.</li></ul></li></ul><h4 id="0b27f890-222d-4e1a-b4d3-a7750da83783" data-toc-id="0b27f890-222d-4e1a-b4d3-a7750da83783" collapsed="false" seolevelmigrated="true">Green Chemistry</h4><ul><li>Application of Knowledge: By understanding fundamental concepts like chemical bonding, molecular geometries (shapes), and intermolecular forces, scientists can design and develop new medicines, molecules, and materials.</li><li>Societal and Environmental Impact: These advancements aim to benefit society while concurrently having a minimal adverse impact on the environment and human health.</li><li>Molecular Recognition: This concept involves the specific interaction between molecules, much like a lock and key. Understanding and utilizing molecular recognition in production methods offers significant green chemistry advantages by enabling highly selective and efficient processes, reducing waste, and minimizing the use of harsh reagents.</li></ul><h4 id="bfa3a37d-c329-4e0a-8555-5eda29cd7e3d" data-toc-id="bfa3a37d-c329-4e0a-8555-5eda29cd7e3d" collapsed="false" seolevelmigrated="true">Chemical Vocabulary (Example: Ammonia,$ NH_3 $)</h4><ul><li>Name: Ammonia</li><li>Chemical Formula:$ NH_3 $</li><li>Lewis Formula: $
 H
 \vert
 H-N:
 \vert
 H
 $$
- Molecular Geometry: Trigonal Pyramidal