Periodic Trends, Electron Configuration, and Atomic Structure: Study Notes

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Periodic Trends and Atomic Structure

Aluminum: Low-Density Atoms and Periodic Properties

Aluminum is a low-density metal with unique properties that can be explained by its atomic structure and position in the periodic table. Understanding periodic properties helps predict the behavior of elements.

Periodic Property: A characteristic of elements that shows a predictable pattern across periods or groups in the periodic table (e.g., atomic radius, ionization energy).
Example: Aluminum's low density is due to its relatively large atomic radius and metallic bonding.

The Periodic Table and the Periodic Law

Organization and Trends

The periodic table arranges elements by increasing atomic number, revealing recurring trends in their properties. The periodic law states that the properties of elements are periodic functions of their atomic numbers.

Groups (Columns): Elements in the same group have similar chemical properties due to similar valence electron configurations.
Periods (Rows): Elements in the same period have the same number of electron shells.
Blocks: The table is divided into s, p, d, and f blocks based on the type of atomic orbital being filled.
Example: Alkali metals (Group 1) are highly reactive due to a single valence electron.

Electron Configuration and Atomic Orbitals

Quantum Numbers and Orbitals

Electrons in atoms occupy orbitals defined by quantum numbers. The arrangement of electrons is described by electron configurations, which follow specific rules.

Principal Quantum Number (n): Indicates the main energy level (shell).
Angular Momentum Quantum Number (l): Defines the shape of the orbital (s, p, d, f).
Magnetic Quantum Number (ml): Specifies the orientation of the orbital.
Spin Quantum Number (ms): Indicates the spin direction of the electron (+1/2 or -1/2).
Aufbau Principle: Electrons fill the lowest energy orbitals first.
Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.
Hund's Rule: Electrons occupy degenerate orbitals singly before pairing up.
Example: The electron configuration of oxygen is 1s2 2s2 2p4.

Valence Electrons and the Periodic Table

Valence electrons are the outermost electrons and determine an element's chemical reactivity. The number of valence electrons corresponds to the group number for main-group elements.

Example: Carbon (Group 14) has 4 valence electrons.

Electron Configurations and Elemental Properties

Electron Configurations of Main-Group and Transition Elements

Electron configurations explain the chemical and physical properties of elements, including their placement in the periodic table and their typical oxidation states.

Main-Group Elements: Fill s and p orbitals.
Transition Metals: Fill d orbitals, often resulting in variable oxidation states.
Example: Iron (Fe): [Ar] 4s2 3d6

Periodic Trends

Atomic and Ionic Radii

Atomic radius is the distance from the nucleus to the outermost electron shell. Ionic radius refers to the size of an ion. Both show predictable trends across the periodic table.

Atomic Radius: Decreases across a period (left to right) and increases down a group.
Ionic Radius: Cations are smaller than their parent atoms; anions are larger.
Example: Na+ is smaller than Na; Cl- is larger than Cl.

Ionization Energy (IE)

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It generally increases across a period and decreases down a group.

First Ionization Energy: Energy to remove the first electron.
Successive Ionization Energies: Each subsequent electron requires more energy to remove.
Example: Helium has the highest first ionization energy of all elements.

Electron Affinity (EA)

Electron affinity is the energy change when an atom gains an electron. A more negative value indicates a greater tendency to accept an electron.

Trends: EA becomes more negative across a period and less negative down a group.
Example: Chlorine has a more negative electron affinity than sodium.

Effective Nuclear Charge (Zeff)

Effective nuclear charge is the net positive charge experienced by valence electrons, accounting for shielding by inner electrons.

Formula:

Z: Atomic number (number of protons)
S: Number of inner (core) electrons
Trend: Zeff increases across a period, leading to smaller atomic radii and higher ionization energies.

Key Equations

Coulomb's Law: Describes the force between two charged particles.

F: Force between charges
q1, q2: Magnitudes of the charges
r: Distance between charges
\varepsilon_0: Permittivity of free space

Representative Table: Element Classification

The following table summarizes key classification information for selected elements:

Symbol	Name	Atomic #	Grouping	Orbital Block	Class	Group #	Group Name
K	Potassium	19	Main-Group	s-block	Nonmetal	1	Alkali Metal
Fe	Iron	26	Transition	d-block	Metal	8	Transition Metal
Pb	Lead	82	Main-Group	p-block	Metal	14	Post-transition Metal
Cl	Chlorine	17	Main-Group	p-block	Nonmetal	17	Halogen
Pr	Praseodymium	59	Metalloid	f-block	N/A	6	Lanthanide

Additional info: Some entries inferred for completeness.

Summary of Periodic Trends

Atomic Radius: Decreases across a period, increases down a group.
Ionization Energy: Increases across a period, decreases down a group.
Electron Affinity: Becomes more negative across a period, less negative down a group.
Effective Nuclear Charge: Increases across a period.

Practice Problems and Applications

Predict electron configurations and draw orbital diagrams for various elements and ions.
Arrange elements or ions in order of increasing/decreasing atomic radius, ionization energy, or electron affinity.
Identify periodic trends using only the periodic table.
Apply Coulomb's Law to explain trends in atomic and ionic properties.

Additional info: Practice problems in the file reinforce understanding of periodic trends, electron configurations, and atomic structure.