1/41
Vocabulary flashcards covering atomic structure, electron configurations, periodic trends, bonding types, and properties derived from bonding.
Name | Mastery | Learn | Test | Matching | Spaced |
---|
No study sessions yet.
Atom
The basic unit of matter composed of electrons, protons, and neutrons; bonding depends on its structure.
Electron
Negatively charged subatomic particle with mass about 9.11×10^-31 kg; orbits the nucleus.
Proton
Positively charged subatomic particle with mass about 1.67×10^-27 kg; located in the nucleus.
Neutron
Electrically neutral subatomic particle with mass about 1.67×10^-27 kg; located in the nucleus.
Atomic Number (Z)
Number of protons in the nucleus; equal to the number of electrons in a neutral atom.
Atomic Mass Unit (amu)
Defined as 1/12 the mass of a carbon-12 atom.
Atomic Weight
The mass of one mole of atoms (6.022×10^23 atoms) of an element.
Isotopes
Atoms of the same element with the same Z but different neutron numbers (e.g., Carbon-12, Carbon-13, Carbon-14).
Bohr Model
Electrons revolve around the nucleus in discrete shells with quantized energy levels.
Wave-Mechanical Model
Electrons exhibit wave-particle duality and occupy orbitals—regions of probability.
Principal Quantum Number (n)
Defines the shell or energy level (K, L, M… or 1, 2, 3…).
Subsidiary Quantum Number (l)
Defines orbital shape (s, p, d, f).
Magnetic Quantum Number (ml)
Orientation of the orbital in space.
Spin Quantum Number (ms)
Direction of electron spin (+1/2 or −1/2).
Orbital
Region in space where there is a high probability of finding an electron; defined by n, l, ml.
Energy Level / Shell
Group of orbitals with similar energy, defined by the principal quantum number n.
s, p, d, f Orbitals
Types of atomic orbitals with characteristic shapes used to describe electron locations.
Electron Configuration
Distribution of electrons among orbitals, filling lowest energy levels first.
Hydrogen 1s1
Hydrogen electron configuration: 1s1.
Helium 1s2
Helium stable configuration: 1s2.
Carbon 1s2 2s2 2p2
Carbon electron configuration with 4 valence electrons.
Noble Gases
Elements with completely filled outer shells; chemically stable.
Valence Electrons
Electrons in the incomplete outer shell that participate in bonding.
Periodic Table
Organization of elements by electron structure and electronegativity trends.
Metals
Electropositive elements that tend to lose electrons to form cations.
Nonmetals
Electronegative elements that tend to gain electrons to form anions.
Electronegativity
Atom’s tendency to attract electrons in a bond; measured on scales like Pauling’s.
Pauling Scale
Common electronegativity scale (approximately 0.7 to 4.0).
Ionic Bonding
Bonding via transfer of electrons from metals to nonmetals; strong Coulombic attraction; examples: NaCl, MgO, CaF2.
Covalent Bonding
Bonding via sharing of valence electrons between atoms with similar electronegativity; directional bonds.
Metallic Bonding
Bonding in metals with delocalized valence electrons forming an electron cloud around positive ion cores.
Mixed Ionic-Covalent Bonding
Bonds that have both ionic and covalent character; ionic character estimated from electronegativity differences.
Secondary Bonding (Van der Waals Forces)
Weaker dipole-based interactions (permanent, induced, fluctuating) important in polymers and molecular solids.
Hydrogen Bonding
A strong type of dipole-dipole interaction involving hydrogen and a highly electronegative partner.
Bond Energy (Eo)
Energy required to break a bond; higher bond energy indicates greater stability and higher melting point.
Bond Length (r)
Distance between the centers of two bonded atoms at equilibrium.
Melting Temperature (Tm)
Temperature at which a solid becomes a liquid; higher for stronger bonds.
Thermal Expansion (α)
Coefficient describing how a material’s size changes with temperature; weaker bonds expand more.
Primary Bonding Types
Ionic, covalent, and metallic bonds—the main, strong interatomic bonds.
Secondary Bonding
Weaker, dipole-based interactions such as van der Waals forces.
Bonding and Properties
Bond type and strength influence melting point, strength, conductivity, and thermal expansion.
Periodicity of Bonding Trends
Element groups and electronegativity trends in the Periodic Table influence bonding behavior.