Chemical Bonds Unit Review Flashcards

Vocabulary flashcards covering introduction to chemical bonding, ionic bonds, covalent bonds, metallic bonds, molecular geometry (VSEPR theory), and intermolecular forces based on the provided unit study guide.

Intermolecular forces

Forces of attraction that occur between different molecules or substances.

Intramolecular forces

Forces that hold atoms together within a single substance, also known as chemical bonds.

Chemical bond

A force or mutual attraction that holds atoms together in a substance.

Electronegativity

A measure of the ability of an atom in a chemical bond to attract electrons.

Compound

A substance composed of two or more different elements chemically combined in fixed proportions, possessing properties distinct from its constituent elements.

Chemical formula

A shorthand representation using element symbols and subscript numbers to specify the types and numbers of atoms present in a compound.

Octet Rule

The principle stating that atoms gain, lose, or share electrons in order to achieve a full outer energy level of $8$ valence electrons.

Covalent bond

A type of chemical bond formed when nonmetal atoms share electrons to achieve stability.

Ionic bond

A chemical bond formed through the electrostatic attraction between oppositely charged ions created by the transfer of electrons from a metal to a nonmetal.

Crystal lattice

A three-dimensional network of mutually attracted cations and anions in an ionic compound.

Lewis structure

A structural diagram representing an atom or molecule using the atomic symbol for the nucleus and dots for valence electrons.

Cation

A positively charged ion formed when an atom loses one or more electrons, typically occurring in metals.

Anion

A negatively charged ion formed when an atom gains one or more electrons, typically occurring in nonmetals.

Lone pair

A pair of valence electrons in an atom's outer shell that are already paired up and not shared in a chemical bond.

Bonding pair

Unpaired valence electrons that are shared between atoms to form a covalent bond.

Single covalent bond

A covalent bond in which two atoms share $2$ electrons (one bonding pair), represented by a single dash.

Double covalent bond

A covalent bond in which two atoms share $4$ electrons (two bonding pairs), represented by two dashes.

Triple covalent bond

A covalent bond in which two atoms share $6$ electrons (three bonding pairs), represented by three dashes.

Polyatomic ion

A covalently bonded group of two or more atoms that carries an overall positive or negative charge.

Oxidation number

The positive or negative charge of an ion within a compound, written as a superscript to the right of the element symbol.

Rule of Zero Charge

The principle stating that the sum of positive charges from cations and negative charges from anions in an ionic compound must add up to zero.

Binary ionic compound

An ionic compound composed of exactly two different elements that transfer electrons in an ionic bond.

Metallic bond

A chemical attraction holding metal atoms together formed from a shared sea of delocalized valence electrons.

Delocalized electrons

Valence electrons in metallic bonding that are not bound to any single atom or specific bond and are free to move throughout empty orbitals.

Malleability

The physical property of a material, particularly metals, that allows it to be hammered or rolled into thin sheets.

Ductility

The physical property of a material that allows it to be drawn or pulled into thin wires.

Alloy

A solution or mixture of two or more metals (or a metal and another element) combined to produce unique physical properties.

Polar covalent bond

A covalent bond formed between atoms with different electronegativities, resulting in an unequal sharing of electrons and partial charges.

Nonpolar covalent bond

A covalent bond formed when electrons are shared equally between atoms due to an electronegativity difference of zero.

Molecule

A neutral group of atoms held together by covalent bonds that functions as an independent unit.

Diatomic elements

Elements whose natural form consists of two bonded atoms of the same element ($H_2$, $N_2$, $O_2$, $F_2$, $Cl_2$, $Br_2$, $I_2$).

Resonance

A condition occurring in molecules or ions that cannot be correctly represented by a single Lewis structure, requiring multiple contributing structures.

Acid

A distinct type of molecular compound containing hydrogen atoms that produces hydrogen ions ($H^+$) when dissolved in water.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; states that electron pairs around a central atom repel each other and stay as far apart as possible to determine molecular shape.

Electron domain

A region around a central atom occupied by electrons, including bonding pairs, lone pairs, or multiple bonds.

Dipole-dipole force

A short-range attractive force occurring between the oppositely charged ends of polar molecules with permanent dipoles.

Hydrogen bonding

A strong intermolecular force occurring when a hydrogen atom bonded to a highly electronegative atom ($N$, $O$, or $F$) is attracted to an unshared electron pair on an adjacent electronegative atom.

London dispersion forces

Weak, temporary intermolecular attractions resulting from instantaneous dipoles created by the constant motion of electrons in all atoms and molecules.

Viscosity

A fluid's internal resistance to movement or flow.

Boiling point

The amount of energy needed to overcome intermolecular forces of attraction separating liquid particles into gas.