CHEMISTRY GRADE 11 - VOCABULARY FLASHCARDS (From Notes)

Vocabulary flashcards covering key terms from the lecture transcript.

Atom

The fundamental unit of matter that defines an element; smallest unit that retains the element’s properties.

Dalton's atomic theory

Postulates describing atoms as indivisible, identifying conservation laws and definite proportions in compounds.

Modern atomic theory

Theory recognizing subatomic particles, isotopes, and that atoms are composed of nucleus and electrons with changes in theory over time.

Electron

Subatomic particle with negative charge that orbits the nucleus and produces chemical behavior.

Proton

Positively charged subatomic particle in the nucleus; contributes to atomic number.

Neutron

Electrically neutral subatomic particle in the nucleus; contributes to mass but not charge.

Nucleus

Center of the atom containing protons and neutrons; dense core of the atom.

Atomic number (Z)

Number of protons in the nucleus; defines the identity of an element.

Mass number (A)

Total number of protons and neutrons in the nucleus.

Isotope

Atoms of the same element with the same Z but different A (neutron count).

Relative atomic mass (Ar)

Weighted average mass of an element’s atoms, relative to 1/12 of a carbon-12 atom.

Atomic mass unit (amu)

Unit used to express atomic and subatomic masses; 1 amu is 1/12 the mass of a carbon-12 atom.

Electron configuration

Arrangement of electrons in an atom’s orbitals; indicates valence electrons and trends.

Orbital

Region in an atom where there is a high probability of finding an electron.

Subshell

A division of a principal shell (s, p, d, f) with distinct shapes and energies.

Principal quantum number (n)

Main energy level of an electron; indicates size and energy of the shell.

Azimuthal quantum number (l)

Indicates the shape of the orbital (s, p, d, f). Values: 0 to n-1.

Magnetic quantum number (mℓ)

Orientation of the orbital in space; values range from −ℓ to +ℓ.

Spin quantum number (ms)

Intrinsic spin of an electron; values are +1/2 or −1/2.

Aufbau principle

Electrons fill lowest-energy orbitals first before higher ones.

Pauli exclusion principle

No two electrons in an atom can have identical four quantum numbers.

Hund's rule

Electrons fill degenerate orbitals singly before pairing occurs.

Bohr model

Early model of the atom with electrons in circular orbits around a nucleus; explains some spectral lines.

Bohr radius (a0)

The most probable distance of the electron from the nucleus in the ground state of hydrogen.

Rydberg constant (RH)

Constant used in hydrogen-like atom energy level calculations.

Energy level (En)

Possible energies that an electron in a hydrogen-like atom may have.

Emission spectrum

Spectrum consisting of discrete lines produced when electrons emit photons during transitions.

Electromagnetic radiation

Form of energy that travels as waves or particles (photons) through space.

Photon

Quantum of electromagnetic energy; particle of light with energy E=hv.

Planck's constant (h)

Proportionality constant between energy and frequency of radiation.

Speed of light (c)

Constant 3.0 x 10^8 m/s; the speed of all electromagnetic radiation in vacuum.

Wavelength (λ)

Distance between successive crests of a wave; inversely related to frequency.

Frequency (ν)

Number of cycles per second of a wave; measured in Hz.

E = hν

Planck’s relation: energy of a photon equals Planck’s constant times frequency.

Balmer series

Hydrogen emission series with transitions ending at n=2; visible lines.

Orbitals shapes (s, p, d)

s: spherical; p: dumbbell; d: more complex; shapes describe electron density.

Wave-particle duality

Idea that light and matter exhibit both wave-like and particle-like properties.

Heisenberg uncertainty principle

Cannot simultaneously know exact position and momentum of a particle.

Ionic bond

Electrostatic attraction between oppositely charged ions; transfer of electrons often occurs.

Covalent bond

Bond formed by sharing electron pairs between atoms.

Coordinate covalent bond

Covalent bond where both electrons come from the same atom to form a bond.

Lewis symbol

Symbol with valence electrons shown as dots around the element symbol.

Lewis structure

Diagram showing the bonding between atoms and lone pairs of electrons.

VSEPR

Model predicting molecular geometry by repulsion of electron pairs.

Bond angle

Angle between two bonds around a central atom; determined by electron pair repulsion.

Bond length

Distance between nuclei in a chemical bond; related to bond strength.

Hybridization (sp, sp2, sp3)

Mixing of atomic orbitals to form new equivalent hybrid orbitals.

Sigma bond (σ)

Bond formed by end-to-end overlap along the bond axis; single bonds are σ.

Pi bond (π)

Bond formed by sideways overlap of p orbitals; present in double/triple bonds.

Delocalization

Electron density spread over several atoms, as in resonance or metallic bonding.

Metallic bond

Bonding in metals where electrons are delocalized into an electron sea.

Electron sea

Model of metallic bonding in which electrons are free to move around cations.