Electromagnetic Radiation and Atomic Structure
Electromagnetic Radiation (EMR) and Atomic Structure
Electromagnetic Radiation (EMR)
Definition: EMR is the transmission of energy through waves.
Electromagnetic Spectrum: Consists of the following types of EMR:
Gamma rays
X-rays
Ultraviolet
Visible light
Infrared
Microwaves
Radio waves
Properties of Waves
Wavelength (\lambda):
The distance between identical points on successive waves.
Frequency (\nu):
The number of waves that travel through a particular point in one second.
Amplitude:
The vertical distance from the midline of a wave to the top of a peak or the bottom of a trough.
Nature of Light - Speed of a Wave:
The speed of a wave depends on two factors:
The type of wave.
The medium through which the wave is traveling (e.g., air, water).
Speed of Light (c):
In a vacuum, the speed of light (c), frequency (\nu), and wavelength (\lambda) are related through the equation: c = \lambda \nu.
Units: Wavelength is in meters (m), and frequency is in reciprocal seconds (s^{-1}), also known as Hertz (Hz).
It is customary to measure wavelength in units that correspond to their magnitude; for example, visible light is usually measured in nanometers (nm).
Relationships between Wave Properties:
Wavelength and Frequency: Inversely proportional.
Long wavelength implies low frequency.
Energy of a Wave: Directly proportional to its frequency.
High frequency implies high energy.
Atomic Spectra
Photon:
A packet of energy with both particle and wave function.
Photons travel at the speed of light (c).
High-energy photons have short wavelengths and high frequencies.
Atomic Spectra (Line Spectra):
A pattern of colored lines emitted when light from a heated element passes through a prism.
Compare to White Light: White light emits a continuous spectrum (no individual lines).
Uniqueness: Each element has a unique atomic spectrum, producing only certain colors/wavelengths based on its allowed energy levels.
Bohr Model of the Atom
Electron Energy Levels:
Electrons exist in fixed energy levels (called orbits) around the nucleus of an atom.
Energy levels are quantized, meaning electrons can only have certain allowed energies.
The lines of an atomic spectrum are associated with the allowed energies of an element.
Highest energy orbits are furthest from the nucleus.
Ground State:
The lowest possible energy state for an electron.
Excited State:
Energy absorbed by an atom causes an electron to move to a higher energy orbit.
As the energy level (n) increases, the electron becomes more unstable.
Photon Emission (Relaxation):
Energy, in the form of a photon, is emitted when an electron relaxes (transitions) from a high-energy excited state to a lower energy state.
Each transition results in the emission of a photon of a specific energy, corresponding to a line on the atomic spectrum.
Limitations of Bohr's Model:
Cannot accurately describe atoms with more than one electron.
A major change introduced by modern atomic theory is that electrons do not move in fixed orbits.
Modern Atomic Theory
Atomic Orbitals:
Regions in space with a high probability of finding an electron.
Electrons move rapidly within the orbital, creating a high electron density.
Principal Energy Levels (Shells):
Specific regions of space that electrons are confined to.
Numbered as n = 1, 2, 3, \dots
Electrons in the first shell (n=1), closest to the nucleus, are held the most strongly and are very hard to remove.
Electrons in higher numbered shells (further from the nucleus) are held less strongly and are easier to remove.
Energy Levels and Sublevels (Subshells):
Shells are divided into subshells, designated by the letters s, p, d, and f.
Within a given energy level, s subshells have the lowest energy.
The number of subshells in a principal energy level is equal to n.
Orbitals within Subshells:
Within a subshell, electrons are grouped in orbitals.
Orbitals are named by their principal energy level and the subshell type (e.g., 1s, 2p).
The number tells the energy level of the orbital.
The letter tells the type of orbital.
Each type of orbital has a characteristic shape and orientation in space:
s orbital: Spherical and symmetrical.
p orbital: Has a shape similar to a dumbbell.
Electron Spin and Pauli Exclusion Principle
Pauli Exclusion Principle:
States that each orbital can hold a maximum of two electrons.
The two electrons must spin in opposite directions to cancel their magnetic fields (since they have the same charge and would otherwise repel).
Electron Spin:
Electrons generate a magnetic field due to their spin.
Spin is represented by the directions of arrows in an orbital diagram.
Each arrow represents one electron.
Electron Arrangement Rules
Orbital Diagrams:
Used to show the arrangement of electrons for an atom of a given element.
Aufbau Principle:
Electrons fill subshells in order of increasing energy for an atom in the ground state.
Electrons fill the lowest energy subshells first before filling higher energy subshells.
Hund's Rule:
The lowest energy state (most stable) is achieved when a single electron is placed in each orbital of a sublevel first, with each electron having the same spin.
A second electron of opposite spin is then added to each orbital in the sublevel.
Electron Configurations
Definition:
A notation used to express the arrangement of electrons around the nucleus, similar to an orbital diagram.
Noble Gas (Abbreviated) Electron Configurations:
A shorthand method of writing electron configurations.
Uses the fact that noble gases contain the maximum number of electrons in their outermost shell.
Electron configurations of all elements (except H and He) can be written using a noble gas core, represented by the element symbol enclosed in brackets (e.g., [Ar]).
Transition Metals and d-Orbitals:
Transition metals have electrons in d orbitals.
With a few exceptions (e.g., Chromium (Cr), Molybdenum (Mo), Silver (Ag)), d orbitals fill in the expected order according to Hund's rule.
Exceptions occur due to slightly greater stability in having a half-filled or completely filled subshell.
Valence Electrons
Definition:
The outermost electrons, those in the highest energy level.
These are the electrons involved in chemical bonding.
Relationship to Periodic Table:
For elements in the main groups (A groups), the number of valence electrons is equal to the group number.
The period number gives the principal energy level (n) of the valence shell for all elements.
Periodic Properties
Basis of the Periodic Table:
The periodic table works because elements in the same column (group) have the same electron configurations in their outer shell.
Many of an atom's properties correlate with its electron structure and its position in the periodic table.
Atomic Size (Atomic Radius):
Definition: The size of an atom is determined by the size of its outermost electron orbital (e.g., the size of a Chlorine atom is determined by its 3p orbitals).
Trend Down a Group: Atomic size increases down a group because the outermost electrons are farther from the nucleus (higher n).
Trend Across a Period: Atomic size decreases across a period because the outermost electrons are more tightly held due to an increasing number of protons (increased nuclear charge), pulling the electron cloud closer.
Ionization Energy (IE):
Definition: The energy required to remove the most loosely held (outermost) electron from an atom in its gaseous phase.
Example: When Sodium (Na) loses one electron, it becomes a Sodium ion (Na^+). It still has 11 protons but now only 10 electrons, resulting in a +1 charge.
Factors Affecting IE: Electrons (negatively charged) are attracted to the positively charged nucleus.
Electrons with strong attractions to the nucleus will be harder to remove (higher IE).
As an electron gets further from the nucleus, its attraction toward the nucleus decreases (lower IE).
As the number of protons increases, the attraction becomes greater (higher IE).
Trend Down a Group: Ionization energy decreases down a group as the outermost electrons are further from the nucleus (weaker attraction).
Trend Across a Period: Ionization energy increases across a period because the outermost electrons are more tightly held due to an increasing number of protons (stronger attraction).
Metallic Character:
Definition: Metallic elements are those that lose electrons easily.
Trend Down a Group: Metallic character generally increases down a group as the outermost electrons are further from the nucleus and thus easier to remove.
Trend Across a Period: Metallic character decreases across a period because the outermost electrons are more tightly held due to an increasing number of protons, making them harder to lose.