Chemical Bonding and Structure Quiz

Chemical Bonding and Structure

• Chemical bond: A force holding two or more atoms together, creating molecules and crystals.
• Octet rule: During compound formation, atoms gain, lose, or share electrons to achieve a full outer electron shell (eight electrons). This allows atoms to become isoelectronic with noble gases.
• Intramolecular forces: Forces holding atoms together within a molecule.
• Intermolecular forces: Forces between molecules.
• Limitation of octet rule: Some transition and post-transition elements do not follow the octet rule.

Types of Chemical Bonding

• Ionic bonding: Electron transfer from a metal to a non-metal, forming ions. Ionic compounds are formed when metal cations bond with non-metal anions (exception: ammonium ion).
• Covalent bonding: Sharing of electrons between non-metals. Non-polar covalent bonding occurs when electronegativity values are similar, while polar covalent bonding results from significant electronegativity differences.
• Metallic bonding: Delocalized electrons shared among a lattice of positively charged metal ions.

Formation of Ionic Bonding

• Formation of ions involves more than just energy levels but depends on the overall stability compared to elemental forms.
• Lattice energy is a measure of the energy released when ions combine to form a solid ionic compound. It strongly influences melting point, hardness, and solubility of the compound
• Factors affecting ionic bonding include ionization energy and electron affinity.

Lattice Energy

• Lattice energy: Energy required to break one mole of ionic solid into gaseous ions.
• Factors that affect lattice energy include: ion size and magnitude of charge.
• Usually measured in kilojoules per mole (kJ/mol).
• Lattice energy is crucial to predicting properties of ionic compounds.

Factors Affecting Formation of Ionic Bonding

• lonization energy: Elements with lower ionization energy are more likely to form cations.
• Electron affinity: Elements with higher electron affinity are more likely to form anions.
• Lattice energy: Larger lattice energies favor the formation of ionic compounds.
• Exceptions to the octet rule exist in ionic compounds (e.g., less than an octet or more than an octet).

Properties of Ionic Compounds

• Typically crystalline solids with high melting and boiling points.
• Generally soluble in water but insoluble in non-polar solvents,
• Often conduct electricity when molten or dissolved.

Covalent Bonds and Molecular Geometry

• Covalent bonds involve electron sharing.
• Bond length: The optimum distance between atomic nuclei.
• Lone pair: A pair of electrons that are not involved in bonding.
• Molecular geometry is the spatial arrangement of atoms in a molecule.

Lewis Structures

• Simplifies the representation of valence electrons in molecules.
• Electrons are represented as dots or lines.

Single, Double, and Triple Bonds

• Single bonds: Share one pair of valence electrons.
• Double bonds: Share two pairs of valence electrons.
• Triple bonds: Share three pairs of valence electrons.
• Bond length and strength vary with the number of shared electron pairs, with triple bonds being the strongest and shortest.

Coordinate Covalent Bonds

• Bonds in which both electrons are donated from the same atom.

Resonance Structures

• Multiple Lewis structures that contribute to the true electronic structure of a molecule.
• Resonance structures have identical connectivity but different arrangements of electrons.

Exceptions to the Octet Rule in Covalent Bonding

• Less than octet
• More than octet
• Odd number of electrons

Polar and Non-Polar Covalent Molecules

• Polar covalent bonds: Unequal sharing of electrons due to electronegativity differences.
• Non-polar covalent bonds: Equal sharing of electrons due to similar electronegativities

Dipole Moment

• Dipole moment: A measure of the polarity of a molecule.
• Occurs when there's a separation of positive and negative charges or dipoles within the molecule, often in polar covalent bonds
• Dipole moment is measured in debyes (D).

Intermolecular Forces

• Intermolecular forces are forces between molecules
• dipole-dipole forces act between molecules possessing permanent dipoles
• London dispersion forces – atoms form temporary dipoles, temporary dipoles are attractive forces
• Hydrogen bonds are strong dipole-dipole forces involving a hydrogen atom covalently bonded to highly electronegative atoms (especially Nitrogen, oxygen, or fluorine)

Metallic Bonding

• Sharing of delocalized electrons among a lattice of positively charged metal ions.
• Strength depends on the number of valence electrons and the packing of metal atoms.

Chemical Bonding Theories

• Valence bond theory explains covalent bonding through orbital overlap.
• Both Valence shell electron pair repulsion (VSEPR) theory and hybridization are used to predict the shape.
• Molecular orbital theory uses mathematical methods that considers the whole molecule to describe bonding; electron distributions are called molecular orbitals.

Molecular Shape and Molecular Polarity

• Bond polarity and dipole moment
• Overall polarity depends on both bond polarity and molecular shape.

Properties of Covalent Compounds

• Low melting and boiling points due to weak intermolecular forces.
• Poor conductors of heat and electricity due to the lack of free electrons.

Molecular Geometry

• Describes the spatial arrangement of atoms in a molecule.
• Common shapes include linear, trigonal planar, tetrahedral, trigonal pyramidal, and octahedral
• The Valence Shell Electron Pair Repulsion (VSEPR) theory helps predict molecular shapes.

Types of Hybridization

• sp, sp2, sp3, sp3d, sp3d2, sp3d3 (each type uses a mix of s, p, or d orbitals to create new hybrid orbitals, used to predict molecular geometry and shape.

Molecular Orbital Theory (MOT)

• Explains covalent bonding via molecular orbitals.
• Molecular orbitals are formed from combinations of atomic orbitals.
• Bonding and anti-bonding orbitals are distinguished by electron density distribution

Types of Crystals

• Ionic, covalent network, molecular, and metallic,
• These structures affect the properties of solids; bonding type influences the strength, melting point, conductivity, and other physical characteristics of the solid