Forces Between Particles and Chemical Bonding

Core Concepts of Chemical Bonding

• Definition of Chemical Bonding: Chemical bonding refers to the forces that hold adjacent atoms or ions together to form distinct chemical compounds.

• Primary Types of Chemical Bonds:

  • Covalent Bonding: Involves the sharing of electron pairs between valence shells (outermost electron shells) of nonmetal atoms, forming molecular compounds.

  • Ionic Bonding: Involves the complete transfer of valence electrons from one atom (typically a metal) to another (typically a nonmetal), resulting in electrostatic forces that hold the oppositely charged ions together.

  • Metallic Bonding: Involves valence electrons that are delocalized and free to move throughout a lattice of positive metal cations (mobile valence electrons).

• Continuum of Bonding: In most chemical compounds, electrons are shared unequally between atoms. The demarcation line between pure covalent and pure ionic bonding is continuous rather than strictly discrete.

Ionic Bonding and Ion Formation

• Definition of Ionic Bond: An electrostatic force of attraction that holds oppositely charged cations and anions together in an ionic solid.

• Mechanism of Formation:

  • Cation Formation (Oxidation): Metal atoms lose electrons to achieve a stable configuration.

  • Example (Lithium): $\text{Li} \rightarrow \text{Li}^+ + e^-$

  • Anion Formation (Reduction): Nonmetal atoms gain electrons to achieve a stable configuration.

  • Example (Fluorine): $\text{F} + e^- \rightarrow \text{F}^-$

  • Net Ionic Pair Formation: $\text{Li}^+ + \text{F}^- \rightarrow \text{LiF}$

• Predicting Ion Charges (The Octet Rule):

  • Representative elements gain or lose electrons to obtain a completely filled outer shell of 8 valence electrons (an octet), attaining the electron configuration of the nearest noble gas.

• Monatomic Ion Charges by Periodic Group:

  • Group 1A (Alkali Metals): Form $\text{+1}$ cations ($\text{H}^+$, $\text{Li}^+$, $\text{Na}^+$, $\text{K}^+$, $\text{Rb}^+$, $\text{Cs}^+$).

  • Group 2A (Alkaline Earth Metals): Form $\text{+2}$ cations ($\text{Mg}^{2+}$, $\text{Ca}^{2+}$, $\text{Sr}^{2+}$, $\text{Ba}^{2+}$).

  • Group 3A: Form $\text{+3}$ cations ($\text{Al}^{3+}$).

  • Group 5A (Pnictogens): Form $\text{-3}$ anions ($\text{N}^{3-}$ nitride, $\text{P}^{3-}$ phosphide).

  • Group 6A (Chalcogens): Form $\text{-2}$ anions ($\text{O}^{2-}$ oxide, $\text{S}^{2-}$ sulfide, $\text{Se}^{2-}$ selenide, $\text{Te}^{2-}$ telluride).

  • Group 7A (Halogens): Form $\text{-1}$ anions ($\text{H}^-$ hydride, $\text{F}^-$ fluoride, $\text{Cl}^-$ chloride, $\text{Br}^-$ bromide, $\text{I}^-$ iodide).

  • Transition & Post-Transition Metals (Variable & Fixed Charges):

  • Titanium: $\text{Ti}^{4+}$

  • Chromium: $\text{Cr}^{2+}$, $\text{Cr}^{3+}$

  • Manganese: $\text{Mn}^{2+}$

  • Iron: $\text{Fe}^{2+}$, $\text{Fe}^{3+}$

  • Cobalt: $\text{Co}^{2+}$, $\text{Co}^{3+}$

  • Nickel: $\text{Ni}^{2+}$

  • Copper: $\text{Cu}^+$, $\text{Cu}^{2+}$

  • Zinc: $\text{Zn}^{2+}$

  • Silver: $\text{Ag}^+$

  • Cadmium: $\text{Cd}^{2+}$

  • Mercury: $\text{Hg}_2^{2+}$ (mercury(I)), $\text{Hg}^{2+}$ (mercury(II))

  • Tin: $\text{Sn}^{2+}$

  • Lead: $\text{Pb}^{2+}$

  • Bismuth: $\text{Bi}^{3+}$

Naming Ionic Compounds and Polyatomic Ions

• Rules for Naming Binary Ionic Compounds:

  • State the name of the metal (cation) first without modification.

  • State the name of the nonmetal (anion) second, modifying its ending to -ide.

  • Examples:

  • $\text{NaCl}$: Sodium chloride

  • $\text{BaS}$: Barium sulfide

  • $\text{CaF}_2$: Calcium fluoride

• Polyatomic Ions:

  • Definition: Groups of two or more covalently bonded atoms that carry an overall net electrical charge.

• Reference Table of Common Polyatomic Ions:

  • Cation:

  • $\text{NH}_4^+$: Ammonium ion

  • Anions Based on Group 4A:

  • $\text{CN}^-$: Cyanide ion

  • $\text{CH}_3\text{CO}_2^-$: Acetate ion

  • $\text{CO}_3^{2-}$: Carbonate ion

  • $\text{HCO}_3^-$: Hydrogen carbonate ion (bicarbonate ion)

  • $\text{C}_2\text{O}_4^{2-}$: Oxalate ion

  • Anions Based on Group 5A:

  • $\text{NO}_2^-$: Nitrite ion

  • $\text{NO}_3^-$: Nitrate ion

  • $\text{PO}_4^{3-}$: Phosphate ion

  • $\text{HPO}_4^{2-}$: Hydrogen phosphate ion

  • $\text{H}_2\text{PO}_4^-$: Dihydrogen phosphate ion

  • Anions Based on Group 6A:

  • $\text{OH}^-$: Hydroxide ion

  • $\text{SO}_3^{2-}$: Sulfite ion

  • $\text{SO}_4^{2-}$: Sulfate ion

  • $\text{HSO}_4^-$: Hydrogen sulfate ion (bisulfate ion)

  • Anions Based on Group 7A:

  • $\text{ClO}^-$: Hypochlorite ion

  • $\text{ClO}_2^-$: Chlorite ion

  • $\text{ClO}_3^-$: Chlorate ion

  • $\text{ClO}_4^-$: Perchlorate ion

  • Anions Based on Transition Metals:

  • $\text{CrO}_4^{2-}$: Chromate ion

  • $\text{Cr}_2\text{O}_7^{2-}$: Dichromate ion

  • $\text{MnO}_4^-$: Permanganate ion

• Practice Problems: Writing Formulas for Ionic Compounds:

  • Barium and Sulfate: $\text{Ba}^{2+}$ and $\text{SO}_4^{2-} \rightarrow \text{BaSO}_4$

  • Magnesium and Hydroxide: $\text{Mg}^{2+}$ and $\text{OH}^- \rightarrow \text{Mg(OH)}_2$

  • Calcium and Phosphate: $\text{Ca}^{2+}$ and $\text{PO}_4^{3-} \rightarrow \text{Ca}_3(\text{PO}_4)_2$

  • Aluminum and Sulfur: $\text{Al}^{3+}$ and $\text{S}^{2-} \rightarrow \text{Al}_2\text{S}_3$

  • Sodium and Bromine: $\text{Na}^+$ and $\text{Br}^- \rightarrow \text{NaBr}$

  • Sodium and Oxygen: $\text{Na}^+$ and $\text{O}^{2-} \rightarrow \text{Na}_2\text{O}$

Covalent Bonds and Naming Binary Molecular Compounds

• Characteristics of Covalent Compounds:

  • Formed by the sharing of electron pairs between nonmetal atoms ($s$- and $p$-block elements).

  • The presence of a metal in a formula is a strong indicator that the compound is ionic; nonmetal-nonmetal combinations indicate covalent molecular structures.

• Rules for Naming Binary Covalent Compounds:

  • Place the name of the less electronegative element first.

  • Take the stem of the name of the more electronegative element and add the suffix -ide.

  • Use Greek numerical prefixes to denote the number of atoms of each element present in the molecule.

  • Omit the prefix mono- for the first element in the name.

• Greek Numerical Prefixes:

  • $1$: mono-

  • $2$: di-

  • $3$: tri-

  • $4$: tetra-

  • $5$: penta-

  • $6$: hexa-

  • $7$: hepta-

  • $8$: octa-

  • $9$: nona-

  • $10$: deca-

• Practice Problems: Naming Binary Covalent Compounds:

  • $\text{N}_2\text{O}_5$: Dinitrogen pentoxide

  • $\text{CS}_2$: Carbon disulfide

  • $\text{H}_2\text{O}$: Dihydrogen monoxide (water)

  • $\text{CO}$: Carbon monoxide

• Practice Problems: Ionic vs. Molecular Classification and Naming:

  • $\text{Li}_2\text{O}$: Ionic — Lithium oxide

  • $\text{N}_2\text{Cl}_2$: Molecular — Dinitrogen dichloride

  • $\text{MgCl}_2$: Ionic — Magnesium chloride

  • $\text{CO}_2$: Molecular — Carbon dioxide

Lewis Symbols and Lewis Dot Structures

• Lewis Symbols for Atoms:

  • Represent the atomic nucleus and inner core electrons using the element symbol.

  • Treat the four sides around the symbol as positions (boxes) capable of holding up to two electrons each.

  • Valence shell electrons are drawn as dots.

  • Unpaired dots are placed in each position before pairing them up.

• Lewis Dot Structures for Molecules:

  • Covalent bonds are formed by shared electron pairs, represented either as a pair of dots or as a straight line.

  • Example ($\text{H}_2$ molecule): Two hydrogen atoms ($\text{H}\cdot$) combine to share electrons, forming $\text{H}:\text{H}$ or $\text{H}-\text{H}$.

• Bonding Pairs vs. Lone Pairs:

  • Bonding Pair: Electron pairs involved in covalent bonds between atoms.

  • Lone Pair (Nonbonding Pair): Pairs of valence electrons that belong exclusively to a single atom and do not participate in bonding.

  • Example (Hydrogen Fluoride, $\text{HF}$):

* Contains 1 single bonding pair ($\text{H}-\text{F}$) and 3 lone pairs on the fluorine atom to complete its octet.

• Step-by-Step Guide to Drawing Lewis Structures:

  • Step 1: Calculate the sum of all valence electrons from all constituent atoms.

  • For polyatomic anions: Add $1\,e^-$ for each unit of negative charge.

  • For polyatomic cations: Subtract $1\,e^-$ for each unit of positive charge.

  • Step 2: Draw the fundamental skeletal structure.

  • Place the least electronegative element in the central position.

  • Hydrogen ($\text{H}$) and fluorine ($\text{F}$) must always occupy terminal positions.

  • Step 3: Connect the central atom to each terminal atom using a single covalent bond (one electron pair represented by a line).

  • Step 4: Distribute the remaining electrons as lone pairs (dots) to complete octets for all terminal atoms ($8\,e^-$ for octet, except $\text{H}$ which requires $2\,e^-$). Assign any leftover electrons to the central atom.

  • Step 5: If the central atom lacks a complete octet, convert one or more lone pairs from terminal atoms into double or triple bonds by sharing electrons. Move existing electrons; do not add new electrons.

• Lewis Structure Examples & Practice:

  • $\text{H}_2\text{O}$: Central $\text{O}$ with two single bonds to $\text{H}$ atoms and 2 lone pairs on $\text{O}$.

  • $\text{SF}_2$: Central $\text{S}$ with two single bonds to $\text{F}$ atoms, 2 lone pairs on $\text{S}$, and 3 lone pairs on each $\text{F}$ atom.

  • $\text{CO}_2$: Central $\text{C}$ double bonded to two $\text{O}$ atoms ($\text{O}=\text{C}=\text{O}$), with 2 lone pairs on each $\text{O}$ atom.

  • $\text{N}_2$: Triple bond between two $\text{N}$ atoms ($:\text{N}\equiv\text{N}:$), with 1 lone pair on each $\text{N}$ atom.

  • $\text{OSCl}_2$: Central $\text{S}$ bonded to $\text{O}$ and two $\text{Cl}$ atoms.

  • $\text{CH}_2\text{CHCN}$ (Acrylonitrile): Skeletal arrangements with single, double, and triple bonds satisfying octets.

  • Structure Evaluation:

  1. Formaldehyde ($\text{H}_2\text{CO}$): Valid structure.

  2. Chloramine ($\text{NH}_2\text{Cl}$): Valid structure.

  3. Nitric Acid ($\text{HNO}_3$): Valid structure.

  4. Nitrous Oxide ($\text{N}_2\text{O}$): Evaluating octets shows invalid formal configurations when lone pairs are incorrectly placed.

VSEPR Theory and Molecular Geometry

• Valence-Shell Electron-Pair Repulsion (VSEPR) Theory:

  • Premise: Valence electron pairs surrounding a central atom repel one another due to like charges. They orient themselves in three-dimensional space as far apart as possible to minimize repulsions.

  • Central Atom: Any atom in a molecule or polyatomic ion bonded to two or more other atoms.

• Key VSEPR Principles:

  • All valence-shell electron pairs (bonding domains and lone pairs) surrounding the central atom are counted equally when establishing the spatial electron-pair geometry.

  • Double and triple bonds are treated as a single electron-dense area (single domain) when determining geometry.

• Analogy of Electron-Dense Areas:

• 2 Balloons (2 Domains): Linear geometry ($180^\circ$ bond angle).

• 3 Balloons (3 Domains): Trigonal planar geometry ($120^\circ$ bond angle).

• 4 Balloons (4 Domains): Tetrahedral geometry ($109.5^\circ$ bond angle).

• Comprehensive VSEPR Molecular Geometry Table:

  • 2 Electron-Dense Areas:

  • Electron-Pair Geometry: Linear

  • 0 Lone Pairs: Molecular Geometry = Linear

  • 3 Electron-Dense Areas:

  • Electron-Pair Geometry: Trigonal planar

  • 0 Lone Pairs: Molecular Geometry = Trigonal planar

  • 1 Lone Pair: Molecular Geometry = Bent

  • 4 Electron-Dense Areas:

  • Electron-Pair Geometry: Tetrahedral

  • 0 Lone Pairs: Molecular Geometry = Tetrahedral

  • 1 Lone Pair: Molecular Geometry = Trigonal pyramidal

  • 2 Lone Pairs: Molecular Geometry = Bent

• Practice Problems: Predicting Shapes using VSEPR Theory:

  • $\text{CO}_2$: Central $\text{C}$ has 2 electron-dense areas (2 double bonds) and 0 lone pairs. Shape: Linear.

  • $\text{NH}_3$: Central $\text{N}$ has 4 electron-dense areas (3 single bonds, 1 lone pair). Shape: Trigonal pyramidal.

  • $\text{CH}_4$: Central $\text{C}$ has 4 electron-dense areas (4 single bonds, 0 lone pairs). Shape: Tetrahedral.

  • $\text{NH}_4^+$: Central $\text{N}$ has 4 electron-dense areas (4 single bonds, 0 lone pairs). Shape: Tetrahedral.

Electronegativity, Bond Polarity, and Molecular Polarity

• Definitions:

  • Electronegativity ($\text{EN}$ or $\chi$): The quantitative measure of an atom's ability in a covalent bond to attract shared electrons toward itself.

  • Polar Covalent Bond: A covalent bond in which electrons are shared unequally due to a difference in electronegativity between the bonded atoms.

  • Nonpolar Covalent Bond: A covalent bond in which bonding electrons are shared equally between atoms.

  • Bond Polarization: The shifting of bonding electron density toward the more electronegative atom, producing partial positive ($\delta^+$) and partial negative ($\delta^-$) charges.

• Periodic Trends in Electronegativity:

  • Increases from left to right across a period.

  • Decreases from top to bottom down a group.

• Electronegativity Values of Common Elements:

  • $\text{H} = 2.1$

  • Period 2: $\text{Li} = 1.0$, $\text{Be} = 1.5$, $\text{B} = 2.0$, $\text{C} = 2.5$, $\text{N} = 3.0$, $\text{O} = 3.5$, $\text{F} = 4.0$

  • Period 3: $\text{Na} = 0.9$, $\text{Mg} = 1.2$, $\text{Al} = 1.5$, $\text{Si} = 1.8$, $\text{P} = 2.1$, $\text{S} = 2.5$, $\text{Cl} = 3.0$

  • Period 4: $\text{K} = 0.8$, $\text{Ca} = 1.0$, $\text{Ga} = 1.6$, $\text{Ge} = 1.8$, $\text{As} = 2.0$, $\text{Se} = 2.4$, $\text{Br} = 2.8$

  • Period 5: $\text{Rb} = 0.8$, $\text{Sr} = 1.0$, $\text{In} = 1.7$, $\text{Sn} = 1.8$, $\text{Sb} = 1.9$, $\text{Te} = 2.1$, $\text{I} = 2.5$

  • Period 6: $\text{Cs} = 0.7$, $\text{Ba} = 0.9$

• Bond Classification by Electronegativity Difference ($\Delta\text{EN}$):

  • Nonpolar Covalent: $\Delta\text{EN}$ range from $0.0$ to $0.5$ (e.g., $\text{H}_2$ where $\Delta\chi = 0$, $\text{Cl}_2$).

  • Polar Covalent: $\Delta\text{EN}$ range from $0.5$ to $1.7$ (e.g., $\text{HF}$ where $\Delta\chi = 1.8$, $\text{HCl}$).

  • Ionic: $\Delta\text{EN}$ greater than $1.7$ up to $3.3$ (e.g., $\text{LiF}$ where $\Delta\chi = 3.0$, $\text{NaCl}$).

• Polar vs. Nonpolar Molecules:

  • Polar Molecule: A molecule with polar bonds arranged nonsymmetrically, resulting in a net dipole moment ($\mu > 0\,\text{D}$).

  • Nonpolar Molecule: A molecule containing no polar bonds OR containing polar bonds arranged symmetrically so that individual bond dipoles cancel out ($\mu = 0\,\text{D}$).

• Examples of Dipole Moments and Molecular Geometry:

  • Water ($\text{H}_2\text{O}$): Bent geometry, net dipole moment $\mu = 1.85\,\text{D}$. Polar molecule.

• Boron Trifluoride ($\text{BF}_3$): Trigonal planar, symmetric charge distribution. Net dipole moment $\mu = 0\,\text{D}$. Nonpolar molecule.

• Phosgene ($\text{Cl}_2\text{CO}$): Nonsymmetric, net dipole moment $\mu = 1.17\,\text{D}$. Polar molecule.

• Ammonia ($\text{NH}_3$): Trigonal pyramidal, net dipole moment $\mu = 1.47\,\text{D}$. Polar molecule.

• Molecular Polarization Reference Summary Table:

  • $\text{H}_2$: Structural $\text{H}-\text{H}$, No charges, Symmetric distribution -> Nonpolar

  • $\text{HCl}$: Structural $\text{H}-\text{Cl}$, $\delta^+ \text{H}-\text{Cl} \delta^-$, Nonsymmetric charges -> Polar

  • $\text{CO}_2$: Structural $\text{O}=\text{C}=\text{O}$, $\delta^- \text{O} \Leftarrow \text{C}^{2+} \Rightarrow \text{O} \delta^-$, Symmetric distribution -> Nonpolar

  • $\text{BF}_3$: Structural Trigonal Planar, $\text{B}^{3+}$ with three $\delta^- \text{F}$, Symmetric distribution -> Nonpolar

  • $\text{N}_2\text{O}$: Structural $\text{N}\equiv\text{N}-\text{O}$, $\text{N}\equiv\text{N}^+ \Rightarrow \text{O}^-$, Nonsymmetric charges -> Polar

  • $\text{H}_2\text{O}$: Structural Bent, $\delta^+ \text{H} \Rightarrow \text{O}^{2-} \Leftarrow \text{H} \delta^+$, Nonsymmetric charges -> Polar

Interparticle Forces and Physical Properties

• Types of Interparticle Forces:

  • Intramolecular / Primary Bonds: Hold atoms together within a molecule or network (Ionic bonds, Covalent bonds, Metallic bonds).

  • Intermolecular / Secondary Forces: Forces that hold discrete particles (molecules or atoms) together in bulk states of matter.

  • Dipolar Forces (Dipole-Dipole): Electrostatic attraction between the positive end of one polar molecule and the negative end of an adjacent polar molecule.

  • Hydrogen Bonding: Particularly strong dipolar attraction occurring between molecules in which a hydrogen atom is directly bonded to a highly electronegative, small atom ($\text{O}$, $\text{N}$, or $\text{F}$).

  • Dispersion Forces (London Dispersion): Very weak attractive forces present between all atoms and molecules. They arise from momentary, non-symmetric electron cloud distributions creating instantaneous induced dipoles.

• Relative Strengths of Interparticle Forces:

  • Dispersion forces (Weakest, e.g., Nitrogen $\text{N}_2$)

  • Dipolar forces (Intermediate, e.g., Carbon monoxide $\text{CO}$)

  • Hydrogen bonds (Moderate/Strong intermolecular, e.g., Water $\text{H}_2\text{O}$)

  • Ionic bonds (Very strong intramolecular, e.g., Sodium chloride $\text{NaCl}$)

  • Covalent bonds (Very strong intramolecular, e.g., Diamond $\text{C}$)

  • Metallic bonds (Very strong intramolecular, e.g., Iron $\text{Fe}$)

• Impact on Physical Properties:

  • The type and strength of interparticle forces directly influence macroscopic physical properties, including melting points, boiling points, and solubility.

• Practice Problems: Comparing Dispersion Force Strengths in Solids:

  • $\text{K}$ vs. $\text{As}$: $\text{As}$ has stronger dispersion forces due to higher mass and greater polarizability.

  • $\text{F}_2$ vs. $\text{I}_2$: $\text{I}_2$ has a much larger electron cloud and higher polarizability, giving it significantly stronger dispersion forces (solid at room temperature vs. gas for $\text{F}_2$).

  • Pentane ($\text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_2\text{CH}_3$) vs. Octane ($\text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_2\text{CH}_2\text{CH}_2\text{CH}_2\text{CH}_3$): Octane has a longer chain length and larger surface area, resulting in stronger dispersion forces.

• Practice Problems: Identifying All Interparticle Forces:

  • Ethanol ($\text{CH}_3\text{CH}_2\text{OH}$): Dispersion forces, Dipolar forces, and Hydrogen bonding (due to the terminal $\text{-OH}$ group).

  • Carbon Dioxide ($\text{CO}_2$): Nonpolar molecule; exhibits Dispersion forces only.

  • Acetone ($\text{CH}_3\text{COCH}_3$): Contains a polar carbonyl group ($\text{C}=\text{O}$) without direct $\text{O}-\text{H}$ bonds; exhibits Dispersion forces and Dipolar forces.