Glossary
Alloys
Mixtures composed of two or more elements, at least one of which is a metal, created to enhance specific properties.
Example:
Brass, an alloy of copper and zinc, is harder and more durable than pure copper.
Anion
A negatively charged ion, formed when an atom gains one or more electrons.
Example:
A chlorine atom gains an electron to become a chloride anion, Cl⁻.
Cation
A positively charged ion, formed when an atom loses one or more electrons.
Example:
When a sodium atom loses an electron, it becomes a sodium cation, Na⁺.
Coulomb's law
A law describing the electrostatic force between two charged particles, stating that the force is directly proportional to the product of their charges and inversely proportional to the square of the distance between them.
Example:
Coulomb's law explains why ions with higher charges or smaller radii have stronger attractions in an ionic crystal.
Covalent Bonds
Chemical bonds formed when two atoms share one or more pairs of electrons, typically occurring between two nonmetals.
Example:
In a water molecule (H₂O), oxygen and hydrogen atoms are held together by covalent bonds where they share electrons.
Crystal lattice
A highly ordered, three-dimensional arrangement of atoms, ions, or molecules in a crystalline solid.
Example:
The repeating cubic structure of NaCl is an example of a crystal lattice formed by alternating Na⁺ and Cl⁻ ions.
Delocalization
The spreading out of electrons over multiple atoms or bonds, rather than being confined to a single atom or bond.
Example:
In benzene, the pi electrons are delocalized over the entire ring structure, contributing to its stability.
Electrostatic forces
Attractive or repulsive forces between electrically charged particles. In ionic compounds, these strong forces hold cations and anions together.
Example:
The strong electrostatic forces between Na⁺ and Cl⁻ ions are what make table salt a solid at room temperature.
Formal charge
The hypothetical charge an atom would have in a molecule or ion if all electrons in covalent bonds were shared equally between the atoms.
Example:
Calculating the formal charge on atoms in different Lewis structures for the thiocyanate ion (SCN⁻) helps determine the most stable arrangement.
Hybridization
The concept of mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds.
Example:
The carbon atom in methane undergoes sp³ hybridization to form four equivalent bonds with hydrogen atoms, resulting in a tetrahedral shape.
Interstitial Alloys
Alloys formed when smaller atoms fit into the spaces (interstices) between the larger atoms in a metal crystal lattice.
Example:
Steel is an interstitial alloy where small carbon atoms occupy the gaps within the iron crystal structure, making it stronger.
Intramolecular forces
The strong attractive forces that exist *within* a molecule, holding the atoms together to form the molecule itself.
Example:
The intramolecular forces in a molecule of methane (CH₄) are the strong carbon-hydrogen covalent bonds.
Ionic Bonds
Chemical bonds formed between a metal and a nonmetal through the complete transfer of electrons, resulting in the formation of oppositely charged ions.
Example:
Table salt, NaCl, forms through an ionic bond where sodium transfers an electron to chlorine, creating Na⁺ and Cl⁻ ions.
Lattice energy
The energy released when one mole of an ionic compound is formed from its gaseous ions, indicating the strength of the electrostatic forces in the crystal lattice.
Example:
The high lattice energy of magnesium oxide (MgO) compared to sodium chloride (NaCl) is due to the greater charges of Mg²⁺ and O²⁻ ions.
Lewis structures
Diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule.
Example:
Drawing the Lewis structure for water (H₂O) helps visualize the two single bonds and two lone pairs on the oxygen atom.
Metal
An element that tends to lose electrons to form positive ions (cations) and typically exhibits properties like high electrical conductivity, malleability, and ductility.
Example:
Copper is a common metal used in electrical wiring because of its excellent conductivity.
Metallic bonds
A type of chemical bond found in metals, characterized by a 'sea' of delocalized valence electrons shared among a lattice of positively charged metal ions.
Example:
The excellent electrical conductivity of copper is due to its metallic bonds, allowing electrons to move freely.
Molecular Geometry
The three-dimensional arrangement of the atoms in a molecule, determined by the repulsion between electron pairs around the central atom.
Example:
The molecular geometry of ammonia (NH₃) is trigonal pyramidal due to the presence of one lone pair and three bonding pairs around the nitrogen atom.
Molecules
A group of two or more atoms held together by chemical bonds, representing the smallest unit of a chemical compound that can exist independently.
Example:
A single molecule of water consists of two hydrogen atoms bonded to one oxygen atom.
Nonmetal
An element that tends to gain or share electrons to form negative ions (anions) or covalent bonds, and generally lacks metallic properties.
Example:
Oxygen is a nonmetal that readily forms covalent bonds with other nonmetals, like in water (H₂O).
Nonpolar covalent bond
A type of covalent bond in which electrons are shared equally between two atoms, typically occurring when the atoms have similar or identical electronegativities.
Example:
The bond between two oxygen atoms in an O₂ molecule is a nonpolar covalent bond because both atoms have the same electronegativity.
Octet rule
A chemical rule stating that atoms tend to gain, lose, or share electrons in order to achieve a full outer electron shell with eight electrons, similar to noble gases.
Example:
In carbon dioxide (CO₂), both carbon and oxygen atoms satisfy the octet rule by forming double bonds.
Polar covalent bond
A type of covalent bond in which electrons are shared unequally between two atoms due to a difference in electronegativity, creating partial positive and negative charges.
Example:
The O-H bond in water is a polar covalent bond because oxygen is more electronegative than hydrogen, pulling the shared electrons closer to itself.
Potential energy
The energy stored in a system due to its position or arrangement. In chemistry, atoms bond to minimize their potential energy, leading to greater stability.
Example:
When two hydrogen atoms are far apart, they have high potential energy; as they approach and form a bond, their potential energy decreases to a minimum.
Resonance
A concept used when a single Lewis structure cannot adequately describe the bonding in a molecule or ion, requiring multiple valid Lewis structures (resonance structures) to represent the true electron distribution.
Example:
The carbonate ion (CO₃²⁻) exhibits resonance, meaning its actual structure is a hybrid of three possible Lewis structures, with electrons delocalized over all oxygen atoms.
Substitutional Alloys
Alloys formed when atoms of one element replace atoms of another element of similar size in the crystal lattice.
Example:
Bronze, an ancient substitutional alloy of copper and tin, was used for tools and weapons due to its increased hardness.
VSEPR (Valence Shell Electron Pair Repulsion) theory
A model used to predict the 3D geometry of molecules based on the repulsion between electron pairs (both bonding and nonbonding) around a central atom.
Example:
Using VSEPR theory, we can predict that methane (CH₄) has a tetrahedral molecular geometry because its four electron domains repel each other equally.