Atomic Structure and Quantum Theory

Flashcards covering key vocabulary from the 'Atomic Structure' lecture, including definitions of electromagnetic radiation, quantum numbers, electron configurations, and periodic trends.

Electromagnetic Radiation

One of the means by which energy travels through space, exhibiting wavelike behavior and traveling at the speed of light in a vacuum.

Wavelength (!)

The distance between two consecutive peaks or troughs in a wave.

Frequency (")

The number of waves per second that pass through a given point in space.

Amplitude

The height of a wave crest or depth of a trough.

Speed of Light (c)

The constant speed at which electromagnetic radiation travels in a vacuum, approximately 2.9979 x 10^8 m/s.

Classical Theory (Matter)

Proposed that matter is particulate and electrons are particles.

Classical Theory (Energy)

Proposed that energy is continuous and wavelike, and can be absorbed in any amount.

Blackbody Radiation

The characteristic changes in intensity and wavelength of emitted light as an object is heated, which could not be described by classical theory.

Planck's Constant (h)

A proportionality constant (6.626 x 10^-34 J•s) relating the energy of a quantum of electromagnetic radiation to its frequency, stating energy can be gained or lost only in whole-number multiples of h".

Quanta

Units of energy (h"), representing the smallest discrete amount of energy that can be absorbed or emitted by an atom.

Photoelectric Effect

The phenomenon where electrons are emitted from a metal surface when light shines on it, with specific observations regarding threshold frequency and intensity.

Dual Nature of Light

The concept that electromagnetic radiation exhibits both wave-like and particle-like (photon) behavior.

Photons

Stream of particles representing electromagnetic radiation, each carrying a quantum of energy.

Diffraction

The spread of waves (such as X-rays or electrons) as they pass through an aperture or around an obstacle, forming a pattern that demonstrates wavelike behavior.

De Broglie's Equation

An equation (! = h/mv) that relates the wavelength of a particle to its mass and velocity, supporting the idea of wave-particle duality for matter.

Continuous Spectrum

A spectrum of light containing all wavelengths within a given range, such as that produced when white light passes through a prism.

Line Spectrum

A spectrum showing only specific, discrete wavelengths or 'lines' of light, indicative of quantized energy levels within atoms.

Quantized Electron Energy Levels

The concept that electrons in an atom can only occupy certain discrete energy states, not just any energy level.

Bohr's Quantum Model

A model for the hydrogen atom stating that electrons exist in specific allowable energy levels (stationary states) and change states by absorbing or emitting photons.

Quantum Mechanical Model of the Atom

A model where an electron bound to the nucleus is viewed as a standing wave, which remains constrained within a region of space.

Schrödinger's Equation

A mathematical equation (Ĥψ = Eψ) whose solutions (wave functions or orbitals) describe the probability of finding an electron in a particular region of space and its corresponding energy.

Orbital

A specific wave function from Schrödinger's equation, representing a region of space around the nucleus where an electron is most likely to be found.

Heisenberg Uncertainty Principle

A fundamental limitation stating that it is impossible to precisely know both the position and momentum of a particle at the same time.

Electron Probability (ψ²)

Represented by a probability distribution or electron density map, indicating the likelihood of finding an electron at a given point in space.

Principal Quantum Number (n)

A positive integer (1, 2, 3…) that indicates the relative size and energy level of an orbital.

Angular Momentum Quantum Number (l)

An integer from 0 to n-1 that describes the shape of an atomic orbital (l=0 for s, l=1 for p, l=2 for d, l=3 for f).

Magnetic Quantum Number (ml)

An integer from -l to l that describes the orientation of an orbital in space around the nucleus.

s orbitals

Orbitals characterized by their spherical shape, increasing in size with higher n values and having n-1 nodes.

Nodes (Atomic Orbitals)

Areas of zero probability density for finding an electron within an orbital.

p orbitals

Orbitals that have two lobes separated by a node at the nucleus, oriented along the x, y, or z axes.

d orbitals

Orbitals that first appear at the principal quantum level n=3 and have two different fundamental shapes (e.g., cloverleaf).

f orbitals

Orbitals that first occur at the principal quantum level n=4 and have complex shapes.

Degenerate Orbitals

Orbitals that have the same energy level.

Ground State (Atom)

The lowest possible energy state of an atom where its electrons occupy the lowest available orbitals.

Excited State (Atom)

An energy state of an atom higher than its ground state, typically achieved when an electron is transferred to a higher-energy orbital.

Pauli Exclusion Principle

A principle stating that no two electrons in a given atom can have the same set of four quantum numbers, meaning an orbital can hold only two electrons, and they must have opposite spins.

Electron Spin Quantum Number (ms)

A quantum number that can be +½ or –½, referring to the two possible intrinsic spin orientations of an electron.

Polyelectronic Atoms

Atoms with more than one electron, where electron repulsions and attractions cannot be calculated exactly.

Screening/Shielding (Electron Interactions)

The phenomenon where the attraction of an electron to the nucleus is weakened due to repulsions from other electrons (inner electrons) between the electron of interest and the nucleus.

Penetration Effect

The observation that electrons prefer to fill orbitals in the order s, p, d, and then f due to differences in their energy levels, influenced by how effectively they penetrate the electron core.

Electron Configuration

The distribution of electrons of an atom or ion in its atomic orbitals.

Aufbau Principle

A rule stating that electrons fill available orbitals of the lowest possible energy first, building up the electron configuration of an atom.

Hund's Rule

A rule stating that for degenerate orbitals, the lowest energy configuration is achieved by having the maximum number of unpaired electrons with parallel spins.

Inner or Core Electrons

Those electrons an atom has in common with the previous noble gas and any completed transition series; not typically involved in bonding.

Valence Electrons

Electrons involved in bonding; for main group elements, these are all electrons at the highest principal quantum level; for transition elements, ns and (n-1)d electrons are valence.

Isoelectronic Species

Atoms or ions that have the same number and configuration of electrons as another species.

Paramagnetism

A magnetic property exhibited by a species with one or more unpaired electrons, causing it to be attracted by a magnetic field.

Diamagnetism

A magnetic property exhibited by a species with all its electrons paired, causing it to be slightly repelled by a magnetic field.

Ionization Energy

The energy required to remove an electron from a gaseous atom or ion.

Ionization Energy Trend (Period)

Generally increases across a period due to increasing nuclear charge and stronger attraction for electrons.

Ionization Energy Trend (Group)

Generally decreases down a group as the principal quantum number (n) increases, making electrons farther from the nucleus and easier to remove.

Atomic Radii Trend (Group)

Increases down a group as orbital size increases, leading to larger atoms.

Atomic Radii Trend (Period)

Decreases across a period as the effective nuclear charge (Zeff) increases, drawing valence electrons closer to the nucleus.

Ionic Radii Trend

Within a series, ion size generally decreases with increasing nuclear charge (e.g., 3- > 2- > 1- > 1+ > 2+ > 3+).

Effective Nuclear Charge (Zeff)

The net positive charge experienced by an electron in a polyelectronic atom, which is less than the actual nuclear charge due to shielding by other electrons.