Chemical Bonding and Molecular Structure

Periodic Table of Elements

• Overall themes:

  • The periodic table groups elements by similar properties and atomic number order.

  • It includes both naturally occurring and synthetic elements, plus the lanthanide block.

3.1 TYPES OF CHEMICAL BONDS

• The Ionic Bond

  • Cations and anions achieve stable electron configurations by n s^2 n p^6 (octet) in the ionic framework.

  • Cations and anions attract electrostatically.

  • Charge balance: n(+ ) = n(-).

  • New compounds exhibit new properties due to ionic interactions.

  • In solution, ions dissociate (NaCl example).

  • Crystalline lattice structures form due to ionic bonding.

3.5 THE COVALENT CHEMICAL BOND

• The Covalent Bond

  • Involves sharing electrons (n‘e’).

  • Electron clouds overlap to form a bond.

  • Compounds are generally not as strong as ionic solids.

  • Most covalent compounds are insoluble and poor electrical conductors.

• Bond Length

  • The ideal distance at which attractions and repulsions are balanced, yielding a minimum energy.

  • Example: H–H bond distance.

• Examples of covalent compounds: HCl, hydrocarbons, sugars.

3.1 TYPES OF CHEMICAL BONDS (SUMMARY)

• Formation pathways:

  • Na atom and Cl atom form ionic bond via electron transfer: Na^+ and Cl^- forming NaCl (formula unit).

  • Cl–Cl forms covalent bond (Cl2 molecule) via electron sharing.

• Ionic Bond: electrostatic attraction between ions; results in formula units like NaCl.

• Covalent Bond: electron sharing between atoms; results in discrete molecules like Cl2.

3.3 IONS AND IONIC COMPOUNDS

• Predicting formulas and naming Binary Ionic Compounds

  • Type 1 (metal–nonmetal): NaBr → Sodium bromide; MgBr2 → Magnesium bromide; BaO → Barium oxide; general rule: Mg^{2+} + Br^- → MgBr2.

  • Type 2 (transition metals with variable oxidation states): FeCl3 → Iron(III) chloride; PbO → Lead(II) oxide; FeCl2 → Iron(II) chloride; SnS2 → Tin(IV) sulfide.

3.4 PARTIAL IONIC CHARACTER IN COVALENT BONDS

• Pure (non-polar) vs Polar Covalent Bonds

  • Electronegativity defines how strongly an atom attracts electrons in a bond.

  • Polar covalent bonds involve unequal sharing of electrons; comparison metaphor: pooling with a bully (one side contributes more electrons, the other gains advantage).

3.2 ELECTRONEGATIVITY

• Electronegativity scales indicate how badly atoms want electrons.

  • Elements that attract electrons strongly vs those that tend to lose electrons.

• Electronegativity Difference (ΔEN) and Bond Polarity

  • ΔEN helps decide whether a bond is ionic, polar covalent, or nonpolar covalent.

  • Example ranking by increasing ionic character: O-F < C-F < H-F < Na-F

  • Analogy: If strengths are balanced (equal), covalent; if one side dominates, ionic.

3.7 LEWIS SYMBOLS AND STRUCTURES

• Illustrating valence electrons with Lewis symbols

  • Cation and anion examples are shown to depict octet configurations.

  • Octet Rule: Fill valence shells to eight (s^2 p^6).

  • Hydrogen follows the Duet Rule: Fill to two electrons (s^2 equivalent for H).

  • Transition metals have exceptions to the octet rule.

• A Bond is a Shared Electron Pair; A Lone Pair is not shared.

• Double Bond and Triple Bond concepts.

3.7 LEWIS SYMBOLS AND STRUCTURES (Continued)

• Ionic compounds can be represented as Lewis structures with charges.

• Step-by-step approach to simple Lewis structures (Case 1) and more systematic approach (Case 2):

  • Step 1: Count valence electrons; for cations, subtract electrons; for anions, add electrons.

  • Step 2: Build a skeleton with a central atom and single bonds.

  • Step 3: Distribute remaining electrons as lone pairs to satisfy octets (except H).

  • Step 4: Place remaining electrons on central atom as needed.

  • Step 5: Use multiple bonds to satisfy octet where possible.

• Examples: SiH4, CHO2^-, NO^+, OF2.

3.8 EXCEPTIONS

• Odd-Electron Molecules: Achieve octet for the most electronegative atom first.

• Electron-Deficient Molecules and Hypervalent Molecules: Not possible for 2nd period or lower; possible for 3rd period and beyond.

3.9 FORMAL CHARGE AND RESONANCE

• Purpose: Determine the most appropriate Lewis structure.

• Examples show calculating formal charges for different resonance forms.

• Check rules for choosing best structure:
  1) All formal charges zero is preferred if possible.
  2) If nonzero, choose the structure with the lowest magnitude of formal charges.
  3) Prefer adjacent formal charges that are opposite and equal in magnitude.
  4) Negative formal charges should reside on the most electronegative atoms.

• Sum of formal charges equals the overall charge of the species.

3.10 NAMING SIMPLE COMPOUNDS

• Predicting formulas and naming binary covalent compounds:

  • CO → Carbon monoxide; CO2 → Carbon dioxide; SF6 → Sulfur hexafluoride; N2O3 → Dinitrogen trioxide.

• Acids (no oxygen): HCl → Hydrochloric acid; HBr → Hydrobromic acid; HCN → Hydrocyanic acid.

• Acids with oxygen (systematic naming):

  • Suffix -ic added to root if anion ends with -ate (e.g., HNO3 → Nitric acid; H3PO4 → Phosphoric acid).

  • Suffix -ous added if anion ends with -ite (e.g., HNO2 → Nitrous acid; H2SO3 → Sulfurous acid; HClO2 → Chlorous acid).

Polyatomic Ions and Acids (Table)

• Name - Formula - Related Acid - Formula

• ammonium → NH_4^+

• hydronium → H_3O^+

• oxide → O^{2-}

• peroxide → O_2^{2-}

• hydroxide → OH^-

• acetate → CH3COO^- → acetic acid CH3COOH

• cyanide → CN^- → hydrocyanic acid HCN

• azide → N3^- → hydrazoic acid HN3

• carbonate → CO3^{2-} → carbonic acid H2CO_3

• bicarbonate → HCO_3^-

• nitrate → NO3^- → nitric acid HNO3

• nitrite → NO2^- → nitrous acid HNO2

• sulfate → SO4^{2-} → sulfuric acid H2SO_4

• hydrogen sulfate → HSO_4^-

• sulfite → SO3^{2-} → sulfurous acid H2SO_3

• hydrogen sulfite → HSO_3^-

• phosphate → PO4^{3-} → phosphoric acid H3PO_4

• hydrogen phosphate → HPO_4^{2-}

• dihydrogen phosphate → H2PO4^-

• perchlorate → ClO4^- → perchloric acid HClO4

• chlorate → ClO3^- → chloric acid HClO3

• chlorite → ClO2^- → chlorous acid HClO2

• hypochlorite → ClO^- → hypochlorous acid HClO

• chromate → CrO4^{2-} → chromic acid H2CrO_4

• dichromate → Cr2O7^{2-} → dichromic acid H2Cr2O_7

• permanganate → MnO4^- → permanganic acid HMnO4

4.1 MOLECULAR STRUCTURE AND ORBITALS

• Considerations:

  • Bond Angle: the angle between two bonds that include a common atom.

  • Bond Length: the distance between the nuclei of two bonded atoms along the straight line joining the nuclei.

4.1 VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY

• Core idea: Minimize repulsions between electron clouds (bonds or lone pairs).

• Analogy: Imagine balloons; lone pair–lone pair repulsion is the strongest, then lone pair–bonding pair, then bonding pair–bonding pair.

• Relative repulsions: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair; more repulsion for lone pairs and multiple bonds leading to larger distortions.

4.2 MOLECULAR STRUCTURE AND POLARITY

• Electron Pair Geometry and Molecular Shape

  • In-class exercises include predicting geometry for CO2 and BCl3.

  • Follow total electron count, skeletal structure, assign electrons, fill octets, and consider resonance if applicable.

• Example: Water (H2O)

  • Electron pair geometry and molecular shape labeled; Water is bent due to two lone pairs on O.

• Other examples: SF4 (Seesaw geometry), XeF4 (Square planar).

• Multicenter molecules: Some molecules exhibit unconventional shapes (e.g., nitrogen in ammonia-like structures, carbon in CH2, carbon in CO2).

• Local shapes: OH in water, NH in amines, etc.

4.2 MOLECULAR STRUCTURE AND POLARITY

• Molecular Shape vs Electron Pair Geometry

  • Water: electron pair geometry is tetrahedral; molecular shape is bent due to lone pairs.

• Molecular Polarity and Dipole Moment

  • Bond polarity arises from unequal sharing of electrons; dipole moment arises from separation of charges.

  • Overall dipole moment is the vector sum of bond moments.

  • Examples:

  • Nonpolar: symmetrical molecules with zero net dipole moment (e.g., CO2).

  • Polar: asymmetric molecules with net dipole moment (e.g., H2O).

4.3 HYBRID ORBITALS

• 2p and 2s orbitals mixing to form hybrid orbitals

• Hybridization schemes:

  • sp: two equal hybrid orbitals; linear geometry; example: BeH2? (context shows sp usage)

  • sp^2: three equal hybrid orbitals; trigonal planar; example BH3 (borane)

  • sp^3: four equal hybrid orbitals; tetrahedral; example CH4

• Purpose of hybridization: to create energetically equivalent hybrid orbitals for bonding with incoming reactants; reduces mismatch with unhybridized orbitals of different energy.

• Overview sketches:

  • Unhybridized s and p orbitals exist separately before mixing.

  • Hybrid orbitals (sp, sp^2, sp^3) result in equal-energy orbitals for bonding.

σ- and π-bonds

• σ-bond (End-to-end overlap):

  • Formed by end-to-end overlap of orbitals (s–s, s–p, or p–p with end-to-end orientation).

  • Has a node (nodal plane) where probability of finding electrons is zero at the bond center is minimal.

• π-bond (Sideways overlap):

  • Formed by sideways overlap of parallel p-orbitals.

  • Also characterized by nodal planes where electron density is zero.

4.3 MULTIPLE BONDS

• Double Bond:

  • Contains one σ (sigma) bond and one π (pi) bond.

  • Illustration: H–C=C–H with one σ and one π per C=C, plus σ bonds for C–H.

• Triple Bond:

  • Contains one σ bond and two π bonds.

  • Involves overlapping p-orbitals in two perpendicular orientations.

• Example illustrations show σ and π bonds within molecules like C=C, N≡N (in diatomic nitrogen), etc.

4.3 HYBRID ORBITALS (AX geometries)

• Assignment of hybrid orbitals to central atoms follows VSEPR-like ideas (electron-pair geometry) to define the molecular shape:

  • AX2: Linear, 180°, sp

  • AX3: Trigonal Planar, 120°, sp^2

  • AX4: Tetrahedral, 109.5°, sp^3

  • AX5: Trigonal Bipyramidal, 120° and 90°, sp^3d

  • AX6: Octahedral, 90°, sp^3d^2

4.3 HYBRID ORBITALS – ADDITIONAL NOTES

• Practical purpose: hybridization yields energetically equivalent orbitals that better match the symmetry and energy of incoming reactants.

• Visuals show how sp, sp^2, and sp^3 orbitals reorganize the s and p orbitals into new bonding frameworks.

4.3 BONDING AND MOLECULAR POLARITY SUMMARY

• Sigma vs Pi bonds recap:

  • Sigma: end-to-end overlap; single bonds are sigma bonds.

  • Pi: sideways overlap; accompany double and triple bonds.

• Hybridization implications:

  • Determines bond angles and molecular geometry.

  • Explains observed shapes like linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

In-Class Exercises and Key Takeaways

• Lewis structures are used to predict molecular geometry, bond types, and formal charges.

• Practice steps for writing Lewis structures:

  • Count valence electrons; adjust for charges.

  • Create a skeletal structure with central atom and bonds.

  • Distribute remaining electrons as lone pairs to satisfy octets (or duet for H).

  • Use double/triple bonds to satisfy octets where necessary.

• VSEPR predicts electron pair geometry and molecular shape by counting bonding pairs and lone pairs.

• Polarity depends on bond dipoles and molecular geometry; a molecule can have polar bonds but be nonpolar overall if geometry cancels dipoles (e.g., CO2).

• All formal charges should sum to the molecule or ion's overall charge.

• The table of polyatomic ions and acids is essential for quick naming and acid-base chemistry.

• Hybridization concepts connect to observed bond angles and shapes in everyday molecules (e.g., CH4, NH3, H2O).

• The periodic table structure (block organization, lanthanides, actinides, and synthetic elements) is foundational for understanding element properties and bonding behavior.