chemistry chapter 8

Covalent bond

A type of chemical bond where two atoms share electrons.

Molecule

A group of atoms bonded together by covalent bonds.

Sigma (σ) bond

A single covalent bond formed by overlapping orbitals.

Pi (π) bond

A weaker bond found in multiple bonds, caused by side-by-side orbital overlap.

Endothermic reaction

A reaction that absorbs energy (bond-breaking).

Exothermic reaction

A reaction that releases energy (bond-forming).

Single Covalent Bonds

One pair (2 electrons) shared.

Bond length

Distance between nuclei of bonded atoms.

Bond strength

How strongly atoms are held together.

Bond dissociation energy

Energy needed to break a bond.

Octet rule

Atoms form covalent bonds by sharing electrons to achieve a stable configuration of eight valence electrons, similar to noble gases.

Sigma (σ) Bonds vs. Pi (π) Bonds

Sigma bonds are formed by end-to-end orbital overlap and are stronger, while pi bonds are formed by side-by-side orbital overlap and are weaker.

Key Rule for Bond Strength

Shorter bonds = Stronger bonds = More energy needed to break them.

Electrons transferred

Electrons that are completely moved from one atom to another.

Electrons shared

Electrons that are shared between atoms in a covalent bond.

Metal + Nonmetal

A type of bond where electrons are transferred, typically forming ionic compounds.

Nonmetal + Nonmetal

A type of bond where electrons are shared, typically forming molecular compounds.

Bond Strength

The measure of how strongly atoms are bonded together; stronger in ionic bonds and weaker in covalent bonds.

Binary molecular compound

A compound made up of only two nonmetal elements.

Oxyanion

A polyatomic ion that contains oxygen.

Oxyacid

An acid that contains hydrogen, oxygen, and another element.

Binary acid

An acid that contains hydrogen and one other element (no oxygen).

Naming Binary Molecular Compounds

The process of naming compounds made of two nonmetals using prefixes to indicate the number of atoms.

Molecular formula

A formula that shows only the types and number of atoms in a molecule.

Lewis Structure

Diagrams that show the arrangement of atoms, bonding pairs of electrons (shared electrons), and lone pairs of electrons (non-bonding electrons).

Central Atom

The atom with the least attraction for electrons, usually the least electronegative, and hydrogen (H) is never the central atom. The atom most left

Bonding Pairs

Pairs of electrons that are shared between atoms to form covalent bonds.

Resonance

A phenomenon where some molecules cannot be represented by a single Lewis structure and have multiple valid structures.

Odd Number of Electrons

A situation where some molecules have an odd number of valence electrons, such as nitrogen monoxide (NO) which has 11 electrons.

Incomplete Octet

A situation where some elements are stable with fewer than 8 electrons, such as boron trifluoride (BF₃) which has only 6 electrons around boron.

Expanded Octet

A situation where elements in period 3 or higher can have more than 8 electrons, such as phosphorus pentachloride (PCl₅) which has 10 valence electrons.

VSEPR Theory

A model used to predict molecular shapes based on the repulsion between electron pairs around the central atom.

Bond Angle

The angle formed between two terminal atoms and the central atom.

Delocalized Electrons

Electrons that are not associated with a single atom or bond and are spread over several atoms.

Hybrid Structure

The actual structure of a molecule that is an average of all possible resonance structures.

Terminal Atoms

Atoms that are bonded to the central atom in a molecule.

Electron Pair Repulsion

The principle that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion.

Physical Properties

Characteristics of a substance that can be observed without changing its chemical composition, influenced by molecular shape.

Chemical Properties

Characteristics that determine how a substance reacts with other substances, also influenced by molecular shape.

Lone Pairs

Pairs of electrons that are not involved in bonding and are localized on a single atom.

Linear

Two regions of electron density push apart, forming a straight line. 180, sp

Trigonal Planar

Three regions spread out in a flat, triangular shape. 120 sp2

Tetrahedral

Four bonds spread apart evenly in 3D space. 109.5 sp3

Trigonal Pyramidal

One lone pair repels three bonded atoms downward. 107.3 sp3

Bent (V-Shape)

Two lone pairs push the bonded atoms together. 104.5 sp3

Trigonal Bipyramidal

Found in elements with expanded octets. 90,120 sp3d

Octahedral

Found in elements with expanded octets. 90 sp3d2

Hybridization

The process where atomic orbitals mix to form new, equivalent hybrid orbitals.

Electronegativity

The ability of an atom to attract electrons in a chemical bond.

Electronegativity Trends

Increases across a period (left to right) and decreases down a group (top to bottom).

Nonpolar Covalent Bond

A bond where electrons are shared equally between atoms.

Polar Covalent Bond

A bond where electrons are shared unequally, creating partial charges.

Dipole

A molecule with two poles (positive and negative ends).

Dipole Moment

A measure of the separation of positive and negative charges in a molecule.

Molecular Polarity

Determined by bond polarity and molecular shape.

Dispersion Forces (London Forces)

The weakest intermolecular force due to temporary shifts in electron clouds.

Dipole-Dipole Forces

Attraction between polar molecules with permanent dipoles.

Hydrogen Bonding

A strong type of dipole-dipole force occurring between H and F, O, or N.

Covalent Compounds

Compounds formed by the sharing of electrons between nonmetals.

Ionic Compounds

Compounds formed by the transfer of electrons between metals and nonmetals.

Covalent Network Solids

Strong 3D networks formed by covalent bonds, characterized by very high melting points.

Melting/Boiling Point

Affected by the strength of intermolecular forces; high for strong IMFs and low for weak IMFs.

Electrical Conductivity

Covalent compounds are poor conductors unless in water; ionic compounds conduct well in solution or molten.

Solubility

Polar substances dissolve in polar solvents, and nonpolar substances dissolve in nonpolar solvents.

Evaporation Rate

Slow for substances with strong intermolecular forces and fast for those with weak forces.

Key Concept of Polar Molecules

Polar molecules have unequal electron distribution.

Key Concept of Nonpolar Molecules

Nonpolar molecules are symmetrical.