General Chemistry: Atomic Structure, Wave-Particle Duality, and Periodic Trends

Vocabulary practice flashcards covering fundamental chemistry, atomic theory, quantum mechanics, subatomic particles, light phenomena, electron configurations, and periodic trends based on lecture notes.

Matter

Anything that occupies space and has mass.

Substance

A form of matter that has a definite composition and distinct properties.

Chemistry

The study of matter and the changes it undergoes.

Physical Change

A change that does not alter the composition or identity of a substance.

Chemical Change

A change that alters the composition or identity of the substance(s) involved.

Element

A substance that cannot be separated into simpler substances by chemical means; 114 have been identified, 82 occur naturally on Earth, and 32 have been created by scientists.

Compound

A substance composed of atoms of two or more elements chemically united in fixed proportions that can only be separated into pure components by chemical means.

Homogeneous Mixture

A mixture in which the composition is the same throughout.

Heterogeneous Mixture

A mixture in which the composition is not uniform throughout.

Classifications of Matter

The classification scheme dividing matter into mixtures and pure substances, separated by physical or chemical methods.

Extensive Property

A property of a material that depends upon how much matter is being considered, such as mass, length, or volume.

Intensive Property

A property of a material that does not depend upon how much matter is being considered, such as density, temperature, or color.

Dalton's Atomic Theory

Formulated in 1808, it postulates that elements are made of atoms, all atoms of a given element are identical, compounds consist of atoms of multiple elements in integer/simple fraction ratios, and reactions only rearrange atoms.

Law of Conservation of Mass

A principle stating that matter is neither created nor destroyed during a chemical reaction.

Law of Multiple Proportions

The principle that if two elements combine to form more than one compound, the masses of one element combining with a fixed mass of the other are in ratios of small whole numbers.

Cathode Ray Tube

An evacuated tube setup used by J.J. Thomson to deflect cathode rays using magnets and electric fields, allowing him to measure the electron mass-to-charge ratio.

Rutherford's Gold Foil Experiment

An experiment showing that an atom's positive charge is concentrated in the nucleus, protons ($p$) have an opposite charge to electrons ($e^-$), and a proton's mass is $1840 \times$ the mass of an electron ($1.67 \times 10^{-24}\,g$).

Millikan's Oil Drop Experiment

An experiment measuring the fundamental electron charge as $-1.60 \times 10^{-19}\,C$ and determining the electron mass as $9.10 \times 10^{-28}\,g$.

Chadwick's Experiment

A 1932 experiment proving the existence of the neutron ($n$), a neutral particle ($\text{charge} = 0$) in the nucleus with a mass of $1.67 \times 10^{-24}\,g$.

Atomic Number ($Z$)

The number of protons in the nucleus of an atom.

Mass Number ($A$)

The total number of protons and neutrons in the nucleus of an atom ($A = Z + \text{number of neutrons}$).

Isotopes

Atoms of the same element ($X$) that have identical atomic numbers ($Z$) but different numbers of neutrons and mass numbers ($A$).

Dimensional Analysis

A problem-solving method that uses conversion factors to convert a given quantity with certain units to a desired quantity with different units.

Electromagnetic Radiation

The emission and transmission of energy in the form of electromagnetic waves carrying oscillating electric and magnetic field components.

Speed of Light ($c$)

The speed of electromagnetic radiation in a vacuum, equal to $3.00 \times 10^8\,m/s$.

Wavelength ($\lambda$)

The distance between identical points on successive waves.

Amplitude

The vertical distance from the midline of a wave to its peak or trough.

Frequency ($\nu$)

The number of waves that pass through a particular point in 1 second, measured in Hertz ($1\,Hz = 1\,s^{-1}$).

Standing Wave

A stationary wave constrained within a region of space, featuring stationary points called nodes.

Quantization

The limitation of a physical property or system to discrete allowable values rather than a continuous spectrum.

Planck's Constant ($h$)

A physical constant equal to $6.626 \times 10^{-34}\,J\cdot s$ (or $6.63 \times 10^{-34}\,J\cdot s$) used in calculating quantized energy ($E = h\nu$).

Photoelectric Effect

The phenomenon explained by Einstein in 1905 wherein light above a threshold frequency ejects electrons from a metal surface.

Photon

A discrete bundle or particle of light possessing energy proportional to its frequency ($E = h\nu$).

Line Emission Spectrum

A spectrum of discrete light wavelengths emitted when an excited electron transitions from a higher to a lower energy state.

De Broglie Wavelength

The wavelength associated with a particle of mass $m$ moving at velocity $u$, expressed as $\lambda = \frac{h}{m u}$.

Bohr's Model of the Atom

A model proposed in 1913 establishing that electrons occupy quantized circular orbits with discrete energy levels ($E_n = -R_H \frac{Z^2}{n^2}$), valid for hydrogen and hydrogen-like single-electron species.

Heisenberg Uncertainty Principle

The principle stating that it is impossible to simultaneously measure both the position ($\Delta x$) and momentum ($\Delta p$) of a subatomic particle with absolute certainty ($\Delta x \Delta p \ge \frac{h}{4\pi}$).

Schrödinger Wave Equation

An equation describing both wave and particle characteristics of electrons, where a Hamiltonian operator acts on a wave function $\psi(x)$ to equal energy $E$ times $\psi(x)$.

Wave Function ($\psi$)

A mathematical function describing the quantum energy state of an electron, where $\psi^2$ defines the probability density of locating the electron in space.

Pauli Exclusion Principle

A rule stating that no two electrons in an atom can have the exact same set of four quantum numbers ($n, l, m_l, m_s$).

Principal Quantum Number ($n$)

A quantum number ($n = 1, 2, 3, 4, \dots$) defining orbital size and the average distance of an electron from the nucleus.

Angular Momentum Quantum Number ($l$)

A quantum number ($l = 0, 1, 2, \dots, n-1$) indicating the three-dimensional shape of the volume of space an electron occupies ($s, p, d, f$ orbitals).

Magnetic Quantum Number ($m_l$)

A quantum number ($m_l = -l, \dots, 0, \dots, +l$) describing the orientation of an orbital in space.

Spin Quantum Number ($m_s$)

A quantum number taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$ specifying an electron's intrinsic magnetic spin state.

Shell

A collection of atomic orbitals that share the same principal quantum number $n$.

Subshell

A set of orbitals that share both the same principal quantum number $n$ and angular momentum quantum number $l$.

Atomic Orbital

A region of space defined by specific values of $n, l, m_l$ containing up to two electrons with paired spins.

Electron Configuration

The distribution of electrons among atomic orbitals represented as $n l^{\text{number of electrons}}$.

Aufbau Principle

The rule that electrons fill atomic orbitals in order of increasing energy, starting with the lowest available subshell.

Hund's Rule

The rule stating that the most stable electron arrangement in subshells of equal energy is the one with the maximum number of parallel spins.

Valence Electrons

Electrons occupying the orbitals in the outermost principal shell (highest $n$ value) of an atom.

Core Electrons

Inner-shell electrons occupying lower energy levels that match a noble gas electron configuration.

Paramagnetic

A property of substances containing one or more unpaired electrons, causing attraction to an external magnetic field.

Diamagnetic

A property of substances in which all electron spins are paired, causing slight repulsion by a magnetic field.

Paramagnetic vs. Diamagnetic Orbitals

Demonstration of orbital spin states where unpaired electrons yield paramagnetic properties and fully paired electrons yield diamagnetic properties.

Isoelectronic

Having the same total number of electrons and identical ground-state electron configurations.

Transition Metal Cation Rule

When forming cations from transition metals, electrons are always removed first from the $ns$ orbital before being removed from $(n-1)d$ orbitals.

Periodic Law

The principle that physical and chemical properties of elements recur periodically when arranged in order of increasing atomic number.

Periods

The horizontal rows of the periodic table.

Groups

The vertical columns of the periodic table, numbered 1 through 18, containing elements with similar chemical properties.

Effective Nuclear Charge ($Z_{\text{eff}}$)

The net positive nuclear charge felt by an electron in a multi-electron atom, calculated as $Z_{\text{eff}} \approx Z - \sigma$, where $\sigma$ is the core electron shielding constant.

Atomic Radius Trend

The distance metric that increases moving down a group (increasing $n$) and decreases moving left to right across a period (increasing $Z_{\text{eff}}$).

Ionic Radius Rule

Cations are always smaller than the neutral atoms from which they form, whereas anions are always larger than the neutral atoms from which they form.