chapter 5 flashcards

Wave Model of Light

Light is a type of electromagnetic radiation, a form of energy that travels through space as a wave.

Wavelength (λ)

Shortest distance between equivalent points on a continuous wave (e.g., crest to crest). Measured in meters (m), nanometers (nm), etc.

Frequency (ν)

Number of waves passing a point per second. Measured in hertz (Hz), which is equivalent to 1/s.

Amplitude

Height of a wave from the origin to the crest or trough.

Speed of Light (c)

All electromagnetic waves travel at a constant speed in a vacuum.

Wave Equation

Wavelength and frequency are inversely proportional: as one increases, the other decreases.

Electromagnetic Spectrum

Includes all forms of electromagnetic radiation (radio waves, microwaves, infrared, visible light, ultraviolet, X-rays, gamma rays).

Energy and Frequency Relationship

Energy increases with increasing frequency.

Visible Spectrum

Small portion of the spectrum visible to the human eye (red → violet).

Red Light

Long wavelength, low frequency, low energy.

Violet Light

Short wavelength, high frequency, high energy.

Quantized Energy

Max Planck proposed that matter gains/loses energy in specific amounts called quanta.

Energy of a Quantum

Energy is directly proportional to frequency.

Photoelectric Effect

Observed when light of a specific frequency ejects electrons from a metal surface.

Key Findings of Photoelectric Effect

Light below a certain frequency does not eject electrons, regardless of intensity.

Photon

Massless particle carrying a quantum of energy.

Continuous Spectrum

White light produces a full, unbroken spectrum of colors.

Atomic Emission Spectrum

Each element emits specific frequencies of light, forming a unique spectrum (individual lines of color).

Applications of Atomic Emission Spectrum

Identifying elements and analyzing star compositions.

Electromagnetic Radiation

Energy that travels as waves through space.

Planck's Constant (h)

A fundamental constant used in quantum mechanics.

Emission Spectrum

Continuous spectrum includes all wavelengths; emission spectra show only specific wavelengths unique to each element.

Quantum Energy

Matter absorbs/emits energy in whole-number multiples of .

Einstein and Quantum Theory

Used Planck's concept to explain that photons eject electrons only if their energy exceeds a threshold.

Bohr's Model

Electrons move in fixed circular orbits.

Bohr's Model

Orbits correspond to specific energy levels.

Bohr's Model

Only explains hydrogen's atomic emission spectrum accurately.

Quantum Mechanical Model

Electrons treated as waves.

Quantum Mechanical Model

Does not define exact paths (no circular orbits).

Quantum Mechanical Model

Predicts probabilities of electron locations in three-dimensional regions called atomic orbitals.

De Broglie's Wave-Particle Duality

Electrons exhibit wave-like behavior, meaning only certain wavelengths and energies are possible.

Heisenberg Uncertainty Principle

It's impossible to know both an electron's position and velocity simultaneously, leading to the idea of probability clouds instead of fixed paths.

Energy Levels

Energy levels are numbered (n = 1, 2, 3, etc.) and determine the size and energy of orbitals.

Sublevels

Each energy level has sublevels (s, p, d, f), with increasing complexity as n increases.

Atomic Orbitals

Regions where electrons are most likely to be found.

Ground State

The lowest energy state of an atom.

Quantum Number (n)

Describes the size, energy, and shape of atomic orbitals.

de Broglie Equation

Relates a particle's wavelength to its mass and velocity.

Heisenberg Uncertainty Principle

States that the position and velocity of an electron cannot be known simultaneously.

Quantum Mechanical Model

Describes electrons as wave-like entities in probabilistic regions around the nucleus.

Atomic Orbital

A 3D region around the nucleus indicating the probable location of an electron.

Principal Quantum Number

Specifies the energy level of an electron.

Principal Energy Level

Major energy levels in an atom.

Energy Sublevel

Divisions within a principal energy level (s, p, d, f).

Energy States of Hydrogen

Electrons exist in specific energy levels (orbits).

Ground State (n = 1)

The lowest energy level.

Excited State

When energy is added, electrons jump to higher orbits.

Photon Emission

Electrons drop back to lower energy levels by emitting photons, producing specific wavelengths of light.

Hydrogen's Line Spectrum

Electrons transition between energy levels.

Balmer Series

Visible light corresponds to transitions to the second energy level.

Lyman Series

Ultraviolet (drops to n = 1).

Paschen series

Infrared (drops to n = 3).

Limitations of Bohr's model

Explained hydrogen well but failed for multi-electron atoms.

Quantum Mechanical Model

Based on the wave-particle duality of electrons.

Schrödinger

Introduced the quantum mechanical model in 1926.

Electrons as waves

Electrons are treated as waves.

Atomic orbitals

Regions with a high probability of finding an electron.

Energy levels

Quantized, similar to Bohr's model.

de Broglie's Wave-Particle Duality

Proposed that electrons have wave-like properties.

Quantized energy levels

Only specific wavelengths fit within an orbital.

Heisenberg Uncertainty Principle

States that the act of observing an electron changes its position or velocity.

Key implications of the Uncertainty Principle

Impossible to determine an electron's exact location and speed simultaneously.

Schrödinger Wave Equation

Treated electrons as waves, leading to the quantum mechanical model.

Wave Functions

Mathematical solutions describe the probability of finding an electron.

High-density regions

Indicate higher probabilities (electron clouds).

Principal Quantum Number (n)

Indicates energy level and size of orbitals.

Higher n values

Mean larger orbitals and higher energy levels.

Energy Sublevels

Each principal energy level contains sublevels.

Sublevel shapes

s: Spherical, p: Dumbbell-shaped, d and f: More complex shapes.

Electrons in ground and excited states

Figure 10: Electrons in ground and excited states.

Electron transitions and photon emission

Figure 11: Electron transitions and photon emission.

Uneven spacing of hydrogen's energy levels

Figure 12: Uneven spacing of hydrogen's energy levels.

Wave properties of electrons

Figure 13: Wave properties of electrons.

Heisenberg uncertainty principle illustration

Figure 14: Illustration of the Heisenberg uncertainty principle.

Probability cloud of an electron

Figure 15: Probability cloud of an electron around the nucleus.

Essential Questions

How are the Pauli exclusion principle, the aufbau principle, and Hund's rule used to write electron configurations using orbital diagrams and electron configuration notation?

Aufbau principle

Electrons fill orbitals starting with the lowest energy levels.

Pauli exclusion principle

Each orbital can hold a maximum of two electrons with opposite spins.

Hund's rule

In degenerate orbitals (same energy level), one electron is added to each orbital before pairing begins.

Valence electrons

Electrons in the outermost energy level of an atom, responsible for chemical properties.

Electron-dot structures

A shorthand representation of valence electrons as dots around an element's symbol.

Electron configuration

Arrangement of electrons in an atom.

Aufbau Diagram

Shows the relative energy levels of orbitals.

Ground-State Electron Configuration

The most stable arrangement of electrons, where the atom is in its lowest energy state.

Energy Levels and Sublevels

Orbitals fill in order of increasing energy.

Electron Configuration Notation

Format: Principal energy level (e.g., 1, 2, 3...), Sublevel (s, p, d, f), Number of electrons in that sublevel.

Silicon

Orbital Diagram: in 1s, 2s, 2p; Electron Configuration: .

Fluorine

Orbital Diagram: .; Electron Configuration: .

Exceptions to the Aufbau Principle

Chromium: (half-filled d-orbital stability); Copper: (fully filled d-orbital stability).

Example of Valence Electrons

Sulfur () has 6 valence electrons.

Example of Electron-Dot Structure

Oxygen () is represented with its symbol surrounded by dots representing valence electrons.

Rules for Writing Electron Configurations

Start with the lowest energy orbital; Apply the Pauli exclusion principle and Hund's rule to fill orbitals.

Features of the Aufbau Principle

Electrons occupy the lowest-energy orbitals first; Orbitals within the same sublevel are degenerate; Overlap between levels occurs.

Example of Aufbau Principle

Hydrogen (), Lithium ().

Example of Pauli Exclusion Principle

Helium ().

Example of Hund's Rule

Nitrogen (): orbitals are singly occupied before electrons pair.

Summary of Key Points

Electron configurations explain the arrangement of electrons in orbitals; Valence electrons determine chemical properties; Electron-dot structures visually depict valence electrons.