BackElectrons in Atoms and the Periodic Table: Chapter 9 Study Guide
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Electrons in Atoms and the Periodic Table
Blimps, Balloons, and Models of the Atom
The reactivity and inertness of gases such as hydrogen and helium are explained by atomic models. The Hindenburg disaster illustrates the high reactivity of hydrogen, while modern blimps use inert helium for safety. Understanding why hydrogen is reactive and helium is inert is foundational to atomic theory.
Hydrogen is highly reactive due to its electron configuration, making it prone to forming bonds.
Helium is inert because its electron configuration is stable, with a filled shell.
Atomic models help explain these behaviors and the periodic law.
Periodic law: When elements are arranged by increasing atomic number, their properties recur periodically.
Development of Atomic Models
Atomic models evolved to explain the behavior of electrons and the periodic properties of elements. Niels Bohr and Erwin Schrödinger contributed to the development of quantum mechanics, which describes electrons as both particles and waves.
Bohr model: Electrons occupy fixed orbits with quantized energies.
Quantum-mechanical model: Electrons exist in orbitals, which are probability maps rather than fixed paths.
Wave-particle duality: Electrons exhibit both wave-like and particle-like properties.
Light and Electromagnetic Radiation
Nature of Light
Light is a form of electromagnetic radiation, a type of energy that travels through space at a constant speed. It exhibits both wave-like and particle-like behavior, which is fundamental to understanding atomic structure.
Speed of light: m/s
Wavelength (\lambda): Distance between adjacent wave crests.
Frequency (\nu): Number of wave cycles passing a point per second.
Relationship: Wavelength and frequency are inversely related.
Color and the Electromagnetic Spectrum
White light contains a spectrum of wavelengths, which can be separated into colors using a prism. The visible spectrum ranges from red (longest wavelength) to violet (shortest wavelength).
Visible light: 400 nm (violet) to 750 nm (red)
Color perception: Objects appear colored based on the wavelengths they reflect.
Electromagnetic Spectrum
The electromagnetic spectrum includes all types of electromagnetic radiation, from gamma rays (shortest wavelength, highest energy) to radio waves (longest wavelength, lowest energy). Visible light is only a small portion of this spectrum.
Gamma rays: Most energetic, can damage biological molecules.
X-rays: Used in medical imaging, also energetic.
Ultraviolet: Causes sunburn, can damage skin.
Infrared: Felt as heat, emitted by warm objects.
Microwaves: Used in cooking and radar.
Radio waves: Used for communication.
Photons and Energy
Light can be viewed as a stream of particles called photons. The energy of a photon depends on its wavelength: shorter wavelengths carry more energy.
Energy formula: (where is Planck's constant, is frequency)
Violet light: Higher energy per photon than red light.
Atomic Emission Spectra and Models
Atomic Emission Spectra
Each element emits light at specific wavelengths, producing a unique emission spectrum. This is observed in neon signs and lamps filled with different gases.
Emission spectrum: Not continuous; consists of distinct lines.
Hydrogen emission: Contains specific wavelengths corresponding to electron transitions.
The Bohr Model
Niels Bohr developed a model to explain the emission spectra of hydrogen. In this model, electrons occupy quantized orbits, each with a specific energy level.
Quantum number (n): Specifies the energy level of each orbit.
Quantized energy: Electrons can only exist in specific orbits, not between them.
Excitation and Emission
When an atom absorbs energy, an electron is excited to a higher orbit. As it returns to a lower orbit, it emits a photon of light with energy corresponding to the difference between the orbits.
Photon energy:
Emission lines: Each corresponds to a specific electron transition.
Limitations of the Bohr Model
The Bohr model successfully explains hydrogen's emission spectrum but fails for elements with more than one electron. It was replaced by the quantum-mechanical model.
The Quantum-Mechanical Model
The quantum-mechanical model describes electrons as existing in orbitals, which are probability maps rather than fixed paths. This model accounts for the behavior of electrons in all elements.
Orbitals: Regions where electrons are likely to be found.
Probability maps: Show statistical distribution of electron locations.
Principal Quantum Numbers and Orbitals
Orbitals are specified by a principal quantum number (n) and a letter (s, p, d, f) indicating the subshell and shape. The lowest-energy orbital is the 1s orbital.
Ground state: Lowest energy configuration.
Excited state: Electron occupies higher-energy orbital.
Energy increases: With increasing n.
Shapes and Representations of Orbitals
Orbitals can be represented by dot density maps or geometric shapes. The 1s orbital is spherical, with the highest probability of finding the electron near the nucleus.
s orbitals: Spherical shape.
Dot density: Indicates probability of electron location.
Subshells and Principal Shells
The number of subshells in a principal shell equals the value of n. Each subshell is specified by a letter (s, p, d, f).
Shell | Number of Subshells | Letters Specifying Subshells |
|---|---|---|
n = 1 | 1 | s |
n = 2 | 2 | s, p |
n = 3 | 3 | s, p, d |
n = 4 | 4 | s, p, d, f |
Electron Configurations and Orbital Diagrams
Electron configurations show how electrons occupy orbitals. Orbital diagrams use arrows to represent electrons and their spins.
Pauli exclusion principle: Orbitals hold no more than two electrons with opposing spins.
Hund's rule: Electrons fill orbitals singly before pairing.
Valence and Core Electrons
Valence electrons are in the outermost principal shell and are involved in chemical bonding. Core electrons are in inner shells.
Silicon: 4 valence electrons, 10 core electrons.
Selenium: 6 valence electrons, 28 core electrons.
Electron Configurations and the Periodic Table
The periodic table is organized by electron configurations. Elements in the same column have similar valence electron counts and chemical properties.
s block: First two columns.
p block: Last six columns.
d block: Transition metals.
f block: Lanthanides and actinides.
Periodic Trends in Electron Configurations
Main-group elements have valence electrons equal to their group number. Transition metals fill d orbitals, with exceptions for stability.
Group number: Equals number of valence electrons (except helium).
Row number: Equals highest principal quantum number.
Transition metals: d orbital quantum number is row number minus one.
Explanatory Power of the Quantum-Mechanical Model
Noble Gases, Alkali Metals, Alkaline Earth Metals, and Halogens
The quantum-mechanical model explains the stability and reactivity of different groups of elements based on their electron configurations.
Noble gases: Stable due to filled valence shells.
Alkali metals: Highly reactive, tend to lose one electron.
Alkaline earth metals: Lose two electrons to attain noble gas configuration.
Halogens: Gain one electron to attain noble gas configuration.
Periodic Trends
Atomic Size
Atomic size decreases across a period due to increased nuclear charge and increases down a column due to higher principal quantum numbers.
Across a period: Size decreases.
Down a column: Size increases.
Ionization Energy
Ionization energy increases across a period and decreases down a column. It is the energy required to remove an electron from an atom.
Across a period: Ionization energy increases.
Down a column: Ionization energy decreases.
Metallic Character
Metallic character decreases across a period and increases down a column. Metals tend to lose electrons, while nonmetals tend to gain electrons.
Across a period: Metallic character decreases.
Down a column: Metallic character increases.
Summary of Key Concepts
Electromagnetic radiation is energy that travels through space and exhibits both wave and particle properties.
Bohr model explains hydrogen's emission spectrum with quantized orbits.
Quantum-mechanical model describes electron orbitals as probability maps.
Electron configurations show how electrons occupy orbitals, obeying the Pauli exclusion principle and Hund's rule.
Periodic table organizes elements by electron configuration, revealing periodic trends in atomic size, ionization energy, and metallic character.
