🧪AP Chemistry Unit 2 – Molecular and Ionic Bonding

Key Concepts

• Chemical bonds form when atoms share or transfer electrons to achieve a stable electronic configuration
• Ionic bonds involve the complete transfer of electrons from one atom to another, forming oppositely charged ions that attract each other
• Covalent bonds occur when atoms share electrons, creating a stable arrangement between the bonded atoms
• Electronegativity differences between atoms determine the polarity of covalent bonds
  • Nonpolar covalent bonds have equal sharing of electrons (H₂, Cl₂)
  • Polar covalent bonds have unequal sharing of electrons (HCl, H₂O)
• Molecular geometry describes the three-dimensional arrangement of atoms in a molecule, which affects its properties and reactivity
• Intermolecular forces, such as hydrogen bonding and van der Waals forces, govern the interactions between molecules and influence their physical properties

Atomic Structure Basics

• Atoms consist of protons (positively charged), neutrons (neutral), and electrons (negatively charged)
• Protons and neutrons reside in the nucleus, while electrons occupy orbitals surrounding the nucleus
• The number of protons in an atom determines its atomic number and element identity
• Electrons are arranged in energy levels (shells) and sublevels (s, p, d, f) around the nucleus
  • The electron configuration of an atom describes the distribution of electrons in these levels and sublevels (1s², 2s², 2p⁶)
• Valence electrons, those in the outermost energy level, participate in chemical bonding
• Atoms tend to form bonds to achieve a stable octet configuration (8 valence electrons) or duet configuration (2 valence electrons) for elements in the first period

Types of Chemical Bonds

• Chemical bonds are attractive forces that hold atoms together in compounds
• Ionic bonds form between metals and nonmetals through the complete transfer of electrons
  • Metals lose electrons to form positively charged cations
  • Nonmetals gain electrons to form negatively charged anions
• Covalent bonds form between nonmetals through the sharing of electrons
  • Single covalent bonds involve the sharing of one pair of electrons (H₂, HCl)
  • Double and triple covalent bonds involve the sharing of two or three pairs of electrons, respectively (O₂, N₂)
• Metallic bonds occur in metals, where valence electrons are delocalized and shared among all atoms in the metal lattice
• Coordinate covalent bonds (dative bonds) form when one atom donates both electrons in the shared pair (NH₄⁺, H₃O⁺)

Ionic Bonding Explained

• Ionic bonds form when there is a large electronegativity difference between atoms (typically >1.7)
• The more electronegative atom (usually a nonmetal) gains electrons to form an anion, while the less electronegative atom (usually a metal) loses electrons to form a cation
• The resulting oppositely charged ions are held together by strong electrostatic attractions
• Ionic compounds have a crystal lattice structure, where cations and anions are arranged in a repeating pattern
  • The lattice energy is the energy required to separate the ions in the crystal lattice
• Properties of ionic compounds include high melting and boiling points, brittleness, and the ability to conduct electricity when molten or dissolved in water
• Examples of ionic compounds include sodium chloride (NaCl), potassium bromide (KBr), and calcium fluoride (CaF₂)

Covalent Bonding Breakdown

• Covalent bonds form when atoms share electrons to achieve a stable electronic configuration
• The shared electrons are attracted to the nuclei of both atoms, creating a stable bond
• Covalent bonds can be single, double, or triple, depending on the number of electron pairs shared
• Nonpolar covalent bonds occur when electrons are shared equally between atoms with the same electronegativity (H₂, Cl₂)
• Polar covalent bonds occur when electrons are shared unequally due to differences in electronegativity (HCl, H₂O)
  • The more electronegative atom has a partial negative charge (δ-), while the less electronegative atom has a partial positive charge (δ+)
• Covalent compounds have lower melting and boiling points compared to ionic compounds, and they are often soft, malleable, and poor conductors of electricity
• Examples of covalent compounds include water (H₂O), methane (CH₄), and glucose (C₆H₁₂O₆)

Electronegativity and Bond Polarity

• Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond
• The Pauling scale is used to quantify electronegativity, with values ranging from 0.7 (francium) to 4.0 (fluorine)
• The difference in electronegativity between bonded atoms determines the polarity of the bond
  • Nonpolar covalent bonds have an electronegativity difference of 0 to 0.4
  • Polar covalent bonds have an electronegativity difference between 0.4 and 1.7
  • Ionic bonds have an electronegativity difference greater than 1.7
• Bond polarity affects the distribution of charge within a molecule and contributes to its overall polarity
• Dipole moments are used to quantify the polarity of a molecule, with the direction pointing from the positive to the negative end of the molecule
• Examples of polar molecules include water (H₂O), ammonia (NH₃), and hydrogen fluoride (HF), while examples of nonpolar molecules include carbon dioxide (CO₂), methane (CH₄), and dinitrogen (N₂)

Molecular Geometry

• Molecular geometry refers to the three-dimensional arrangement of atoms in a molecule
• The shape of a molecule is determined by the number of electron groups (bonding and lone pairs) around the central atom and the repulsion between them
• The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular geometries based on minimizing the repulsion between electron groups
• The five basic molecular geometries are linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral
  • Linear geometry has two electron groups arranged at a 180° angle (CO₂, HCN)
  • Trigonal planar geometry has three electron groups arranged at 120° angles (BF₃, SO₃)
  • Tetrahedral geometry has four electron groups arranged at 109.5° angles (CH₄, NH₄⁺)
  • Trigonal bipyramidal geometry has five electron groups, with three equatorial and two axial positions (PCl₅, SF₄)
  • Octahedral geometry has six electron groups arranged at 90° angles (SF₆, [Co(NH₃)₆]³⁺)
• The presence of lone pairs on the central atom can lead to distorted geometries, such as bent (H₂O), trigonal pyramidal (NH₃), and see-saw (SF₄)
• Molecular geometry affects the polarity, reactivity, and physical properties of molecules

Intermolecular Forces

• Intermolecular forces are attractive forces between molecules that influence their physical properties and behavior
• The three main types of intermolecular forces are dipole-dipole forces, hydrogen bonding, and London dispersion forces (van der Waals forces)
• Dipole-dipole forces occur between polar molecules, where the partially positive end of one molecule attracts the partially negative end of another
• Hydrogen bonding is a strong type of dipole-dipole interaction that occurs when hydrogen is bonded to a highly electronegative atom (N, O, or F)
  • Hydrogen bonding is responsible for the unique properties of water, such as its high boiling point and surface tension
• London dispersion forces are weak attractions that result from temporary fluctuations in electron distribution, creating instantaneous dipoles
  • These forces are present in all molecules and are the only intermolecular forces in nonpolar molecules
• The strength of intermolecular forces affects properties such as melting and boiling points, viscosity, and surface tension
• Examples of substances with strong hydrogen bonding include water (H₂O) and ammonia (NH₃), while examples of substances dominated by London dispersion forces include methane (CH₄) and carbon dioxide (CO₂)

Real-World Applications

• Understanding chemical bonding is essential for explaining the properties and behavior of materials
• Ionic compounds, such as sodium chloride (NaCl) and potassium chloride (KCl), are used in various applications, including food preservation, deicing, and water softening
• Covalent compounds, like water (H₂O) and glucose (C₆H₁₂O₆), are crucial for life processes, including hydration, energy storage, and cellular respiration
• Hydrogen bonding plays a vital role in the structure and function of biological molecules, such as DNA and proteins
  • The complementary base pairing in DNA (A-T and G-C) is stabilized by hydrogen bonds
  • The secondary structure of proteins (α-helices and β-sheets) is maintained by hydrogen bonding between amino acid residues
• Intermolecular forces influence the design and performance of materials, such as polymers and adhesives
  • Polytetrafluoroethylene (PTFE, Teflon) has low intermolecular forces, making it non-stick and hydrophobic
  • Adhesives rely on strong intermolecular interactions to bond surfaces together
• Understanding chemical bonding helps in the development of new materials with desired properties, such as high-strength alloys, superconductors, and nanomaterials