chemistry chapter 1 - models of the particulate nature of matter

Vocabulary flashcards covering key concepts from particle matter, atomic structure, spectra, electron configurations, and gas laws.

Mass number

Total number of protons and neutrons in the nucleus of an atom.

Atomic number

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

Isotopes

Atoms of the same element (same atomic number) with different numbers of neutrons, hence different mass numbers.

Atomic mass unit (amu)

Unit used to express atomic/molecular masses; 1 amu = 1/12 the mass of carbon-12.

Carbon-12 standard

Carbon-12 (12C) is the reference standard used to define the atomic mass unit.

Mass spectrometry

Analytical technique to measure masses of atoms/molecules and their isotopic abundances.

Emission spectra

Line spectra produced when atoms emit photons as electrons transition to lower energy levels.

Energy level

Quantized electron energy states in an atom; energy levels are arranged from nearest to farthest from the nucleus.

S orbital

Orbital with l=0; holds up to 2 electrons; spherical shape; one per energy level.

P orbital

Orbital with l=1; holds up to 6 electrons; three orbitals (px, py, pz) oriented along x, y, z.

D orbital

Orbital with l=2; holds up to 10 electrons.

F orbital

Orbital with l=3; holds up to 14 electrons.

Aufbau principle

Electrons fill the lowest-energy orbitals first, following a diagonal (Madelung) rule.

Pauli exclusion principle

No two electrons in an atom can have the same set of four quantum numbers; each orbital holds up to 2 electrons with opposite spins.

Hund's rule

For orbitals of equal energy, electrons occupy them singly with parallel spins before pairing.

Madelung rule (diagonal rule)

Guides the order in which orbitals are filled (e.g., 1s, 2s, 2p, 3s, 3p, 4s, 3d, …).

Ionization energy

Energy required to remove one electron from an atom or ion.

Ionization energy trend

Generally increases across a period (left to right) and decreases down a group; harder to remove electrons across a period.

Maximum electrons in energy level n

2n^2 electrons can occupy the nth energy level.

Bohr model

Early atomic model describing electrons in fixed, quantized orbits around the nucleus.

Quantum mechanical model

Modern model where electrons occupy orbitals—probability clouds solved via Schrödinger equation.

s orbital

Shape: spherical; first energy level; one s orbital per energy level (1s, 2s, 3s, …).

p orbitals (px, py, pz)

Dumbbell-shaped orbitals oriented along x, y, z axes; begin at n=2; three orbitals per energy level.

Electron configuration

Arrangement of electrons in orbitals according to Aufbau, Hund’s, and Pauli principles.

E = hf

Photon energy equals Planck’s constant times frequency; E = h f.

Wavelength (λ)

Distance between successive wave crests; used with frequency to describe light.

Frequency (f)

Number of wave cycles per second; measured in Hz.

Hydrogen line emission

Line spectrum from hydrogen due to electron transitions; includes series like Lyman and Balmer.

Atomic spectral lines

Discrete wavelengths emitted or absorbed by atoms corresponding to electron transitions.

Avogadro constant (NA)

NA ≈ 6.02 × 10^23 particles per mole.

Mole

Amount of substance containing NA particles; 1 mole = 6.02 × 10^23 entities.

Relative atomic mass (Ar)

Weighted average mass of an element's isotopes relative to 1/12 the mass of carbon-12.

Relative molecular mass (Mr)

Sum of the relative atomic masses of atoms in a molecule.

Molar mass (M)

Mass of one mole of a substance (g/mol).

Empirical formula

Smallest whole-number ratio of elements in a compound.

Molecular formula

Actual number of each type of atom in a molecule; may be a multiple of the empirical formula.

Mass percent composition

% by mass of an element in a compound: (mass of element in formula / molar mass of compound) × 100%.

Avogadro's Law

At the same temperature and pressure, equal volumes of gases contain equal numbers of moles.

van der Waals forces (London dispersion)

Intermolecular forces causing deviations from ideal gas behavior; more significant for larger, polarizable molecules.

Polar molecules

Molecules with an uneven distribution of electron density, creating a dipole moment.

Non-polar molecules

Molecules with even electron distribution and no permanent dipole moment.

Elastic collisions

Collisions between gas particles that conserve total kinetic energy (ideal gas assumption).

Ideal gas assumptions

Gas particles are point particles, have no volume, exert no intermolecular forces, and collide elastically.

STP (Standard Temperature and Pressure)

Standard conditions for gases; commonly 0°C (273 K) and 1 atm (101.3 kPa).

Real gases vs. ideal gases

Real gases deviate from ideal behavior due to intermolecular forces and finite molecular size.

Molar volume of an ideal gas

Volume occupied by one mole of an ideal gas at STP: 22.4 L.

Dipole moment

Vector quantity indicating molecule polarity due to uneven charge distribution.

Mass number

Total number of protons and neutrons in the nucleus of an atom.

Atomic number

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

Isotopes

Atoms of the same element (same atomic number) with different numbers of neutrons, hence different mass numbers.

Atomic mass unit (amu)

Unit used to express atomic/molecular masses; 1 amu = 1/12 the mass of carbon-12.

Carbon-12 standard

Carbon-12 (12C) is the reference standard used to define the atomic mass unit.

Mass spectrometry

Analytical technique to measure masses of atoms/molecules and their isotopic abundances.

Emission spectra

Line spectra produced when atoms emit photons as electrons transition to lower energy levels.

Energy level

Quantized electron energy states in an atom; energy levels are arranged from nearest to farthest from the nucleus.

S orbital

Orbital with l=0; holds up to 2 electrons; spherical shape; one per energy level.

P orbital

Orbital with l=1; holds up to 6 electrons; three orbitals (px, py, pz) oriented along x, y, z.

D orbital

Orbital with l=2; holds up to 10 electrons.

F orbital

Orbital with l=3; holds up to 14 electrons.

Aufbau principle

Electrons fill the lowest-energy orbitals first, following a diagonal (Madelung) rule.

Pauli exclusion principle

No two electrons in an atom can have the same set of four quantum numbers; each orbital holds up to 2 electrons with opposite spins.

Hund's rule

For orbitals of equal energy, electrons occupy them singly with parallel spins before pairing.

Madelung rule (diagonal rule)

Guides the order in which orbitals are filled (e.g., 1s, 2s, 2p, 3s, 3p, 4s, 3d, …).

Ionization energy

Energy required to remove one electron from an atom or ion.

Ionization energy trend

Generally increases across a period (left to right) and decreases down a group; harder to remove electrons across a period.

Maximum electrons in energy level n

2n^2 electrons can occupy the nth energy level.

Bohr model

Early atomic model describing electrons in fixed, quantized orbits around the nucleus.

Quantum mechanical model

Modern model where electrons occupy orbitals

—probability clouds solved via Schrödinger equation.

s orbital

Shape: spherical; first energy level; one s orbital per energy level (1s, 2s, 3s, …).

p orbitals (px, py, pz)

Dumbbell-shaped orbitals oriented along x, y, z axes; begin at n=2; three orbitals per energy level.

Electron configuration

Arrangement of electrons in orbitals according to Aufbau, Hund’s, and Pauli principles.

E = hf

Photon energy equals Planck’s constant times frequency; E = h f.

Wavelength (\lambda)

Distance between successive wave crests; used with frequency to describe light.

Frequency (f)

Number of wave cycles per second; measured in Hz.

Hydrogen line emission

Line spectrum from hydrogen due to electron transitions; includes series like Lyman and Balmer.

Atomic spectral lines

Discrete wavelengths emitted or absorbed by atoms corresponding to electron transitions.

Avogadro constant (N_A)

N_A \approx 6.02 \times 10^{23} particles per mole.

Mole

Amount of substance containing N_A particles; 1 mole = 6.02 \times 10^{23} entities.

Relative atomic mass (A_r)

Weighted average mass of an element's isotopes relative to 1/12 the mass of carbon-12.

Relative molecular mass (M_r)

Sum of the relative atomic masses of atoms in a molecule.

Molar mass (M)

Mass of one mole of a substance (g/mol).

Empirical formula

Smallest whole-number ratio of elements in a compound.

Molecular formula

Actual number of each type of atom in a molecule; may be a multiple of the empirical formula.

Mass percent composition

% by mass of an element in a compound: (mass of element in formula / molar mass of compound) \times 100%.

Avogadro's Law

At the same temperature and pressure, equal volumes of gases contain equal numbers of moles.

van der Waals forces (London dispersion)

Intermolecular forces causing deviations from ideal gas behavior; more significant for larger, polarizable molecules.

Polar molecules

Molecules with an uneven distribution of electron density, creating a dipole moment.

Non-polar molecules

Molecules with even electron distribution and no permanent dipole moment.

Elastic collisions

Collisions between gas particles that conserve total kinetic energy (ideal gas assumption).

Ideal gas assumptions

Gas particles are point particles, have no volume, exert no intermolecular forces, and collide elastically.

STP (Standard Temperature and Pressure)

Standard conditions for gases; commonly 0°C (273 K) and 1 atm (101.3 kPa).

Real gases vs. ideal gases

Real gases deviate from ideal behavior due to intermolecular forces and finite molecular size.

Molar volume of an ideal gas

Volume occupied by one mole of an ideal gas at STP: 22.4 L.

Dipole moment

Vector quantity indicating molecule polarity due to uneven charge distribution.

Boyle's Law

At constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure (P1V1 = P2V2).

Charles's Law

At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature (V1/T1 = V2/T2).

Gay-Lussac's Law

At constant volume, the pressure of a fixed mass of gas is directly proportional to its absolute temperature (P1/T1 = P2/T2).