Chemistry Chapter 6 Flashcards

Chemical Bonds

• A chemical bond is the electrical attraction between nuclei and valence electrons in different atoms that holds them together.
• Ionic bonding results from the attraction between large numbers of cations and anions.
• Covalent bonding involves the sharing of electron pairs between atoms.

Types of Covalent Bonds

• Nonpolar covalent bonds share electrons equally between bonded atoms, leading to a balanced charge distribution.
• Polar covalent bonds have an unequal sharing of electrons, resulting in an uneven charge distribution.

Molecular Concepts

• A molecule is a neutral cluster of atoms bonded by covalent bonds.
• Molecular compounds consist of molecules as their simplest units.
• A chemical formula denotes the proportional amounts of various elements in a compound using atomic symbols.

Bonding Characteristics

• A bond length is the average distance between two bonded atoms at minimum potential energy.
• Bond energy is the energy needed to break a chemical bond, forming neutral isolated atoms.
• Atoms typically form bonds to achieve an octet, having eight electrons in their outermost energy level.

Electron Representation

• Electron-dot notation displays only valence electrons around the element's symbol using dots.
• Lewis structures represent molecules, showing nuclei, bonding pairs, and unshared electrons.

Types of Chemical Bonds

• A single bond involves sharing one pair of electrons; a double bond shares two pairs; a triple bond shares three pairs.
• Multiple bonds can consist of either double or triple bonds, increasing bond strength and reducing bond length.

Ionic Compounds and Polyatomic Ions

• Ionic compounds consist of positive and negative ions combined to balance charge.
• Polyatomic ions are groups of covalently bonded atoms that carry a charge.

Forces and Properties

• Metallic bonding results from the attraction between metal atoms and a "sea" of electrons, contributing to properties like malleability and ductility.
• Molecular polarity refers to the uneven distribution of charge within a molecule, influencing interactions.
• Intermolecular forces like dipole-dipole forces and hydrogen bonding dictate the physical properties of substances.

Energy and Stability

• Bond formation typically lowers potential energy, contributing to stability; bonded atoms have lower energy than isolated ones.
• Resonance occurs when multiple valid Lewis structures exist for a molecule, indicating delocalized electrons.

Crystalline Structures

• Ionic compounds exist as crystalline solids with ordered atom arrangements in a crystal lattice, minimizing potential energy.
• The lattice energy represents the energy released when forming an ionic compound from gaseous ions.

Electron Interaction and Bonding

• The concept of VSEPR theory explains molecular shape based on repulsion between electron pairs, helping predict molecular geometry.
• Hybridization involves the mixing of atomic orbitals to form new hybrid orbitals that influence bonding characteristics.

Summary of Bonding Types

• The three primary types of bonds are ionic, covalent, and metallic.
• Bonds have varying strengths and characteristics influenced by atomic properties and arrangements.

Additional Bonding Concepts

• In solid states, ions typically cannot move, explaining ionic compounds' brittleness and hardness.
• The energy required to break bonds equals the energy released during formation, illustrating the stability of bonded atoms.