Atomic Structure and Periodicity Study Notes

Definition: The theory that atoms are the fundamental building blocks of matter, reemerged in the early nineteenth century, championed by John Dalton.
- Key points:
1. Each element consists of extremely small particles known as atoms.
2. All atoms of a given element are identical in mass and properties, while atoms of different elements vary.
3. Atoms are not altered into different elements during chemical reactions; they are neither created nor destroyed.
4. Atoms from different elements combine to form compounds, consistently maintaining the same relative number and type of atoms.

Definition: The total mass of substances present after a chemical process is equal to the mass present before the process.
Importance: This law is foundational to Dalton’s atomic theory.
Illustration:
- Diagram depicting reactants (e.g., lead nitrate and sodium chromate solutions) transforming into products (solid lead chromate) without mass loss. This emphasizes conservation during chemical reactions.

Definition: If elements A and B form two different compounds, the masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers.
Example:
- Carbon Oxides:
- Carbon Oxide I: 57.1% oxygen, 42.9% carbon.
- Carbon Oxide II: 72.7% oxygen, 27.3% carbon.

Background: Originally, Dalton considered the atom the smallest particle. Discoveries indicated the atom consists of smaller parts:
- Electrons: Found via cathode rays.
- Radioactivity: Revealed further subatomic particles including protons and neutrons.

Cathode Rays: Streams of negatively charged particles identified by J. J. Thomson in 1897.
- Charge/Mass Ratio: Thomson calculated the charge/mass ratio of electrons as $1.76 imes 10^8$ coulombs/gram.
Millikan Oil-Drop Experiment:
- Determined the electron charge ( $-1.602 imes 10^{-19}$ C) and used it to calculate the electron's mass: $9.109 imes 10^{-31}$ kg.

Definition: The spontaneous emission of high-energy radiation by an atom.
Pioneers: First observed by Henri Becquerel; studied further by Marie and Pierre Curie.
Types of Radiation Identified:
- Alpha (α) particles: Positively charged.
- Beta (β) particles: Negatively charged, similar to electrons.
- Gamma (γ) rays: Uncharged.

Plum Pudding Model: Proposed by Thomson, suggesting a positive sphere with electrons embedded within.
Rutherford's Gold Foil Experiment:
- Shot α particles at gold foil revealing a dense nucleus around which electrons orbit, leading to the nuclear model of the atom.

Basic Particles:
- Protons (+1 charge) and Neutrons (no charge) located in the nucleus.
- Electrons (-1 charge) surrounding the nucleus.
Relative Mass:
- Protons and neutrons have a relative mass close to 1; the electron’s mass is negligible.

Notation:
- Let X be the atomic symbol, A the mass number (A = Z + N), Z the atomic number (number of protons), N the number of neutrons.
Isotopes: Atoms of an element sharing the same number of protons but differing in neutrons.

Isotope	Number of Protons	Number of Neutrons	Total Atomic Mass
$^{11}C$	6	5	11
$^{12}C$	6	6	12
$^{13}C$	6	7	13
$^{14}C$	6	8	14

Definition: A wave characterized by oscillations in electric and magnetic fields.
Key Characteristics:
- Wavelength (λ): Distance between two consecutive peaks or troughs (measured in nanometers, nm).
- Frequency (ν): Number of waves passing a point per second (measured in hertz, s⁻¹).
Relationships:
- $c = <br />\nu imes au$ , where $c$ is the speed of light (~3.00 x 10^8 m/s) and $au$ is the wavelength.

Example: Wavelength of light from fireworks (650 nm).
- Calculate frequency for wavelength $6.50 imes 10^2$ nm.
Exercise 1: What is the wavelength of blue light with frequency of $6.4 imes 10^{14}$ s⁻¹?

Spectrum Types:
- Continuous Spectrum: All wavelengths visible, typically observed when white light passes through a prism.
- Line Spectrum: Discrete wavelengths specific to an atom (e.g., hydrogen).
Significance:
- Electrons in atoms have quantized energies; their transitions correspond to specific wavelengths and energies.

Key Concepts:
- Electrons occupy only quantized orbits.
- Energy changes involve transitions between these discrete orbits, described by: $E = h<br />\nu$ .
- Limitations: The model works for hydrogen but fails with multi-electron systems.

Developed by Erwin Schrödinger: Incorporates the wave-particle duality of matter through wave functions (Ψ), where $ext{Prob. Density} = ext{Ψ}^2$ .
Quantum Numbers:
- Principal quantum number (n): Energy level.
- Angular momentum quantum number (l): Shape of orbital.
- Magnetic quantum number (mₗ): Orbital orientation.

Aufbau Principle: Electrons fill the lowest energy orbitals first.
Hund's Rule: Electrons will singly occupy all degenerate orbitals before pairing.
Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers.

Magnesium (Mg): 12 electrons; Configuration: $1s^22s^22p^63s^2$ .
Argon (Ar): 18 electrons; Configuration: $1s^22s^22p^63s^23p^6$ .
Sulfur (S): $1s^22s^22p^63s^23p^4$ .
Cadmium (Cd): 48 electrons; Configuration: $1s^22s^22p^63s^23p^64s^23d^{10}4p^65s^24d^{10}$ .

Definition: The repetitive pattern of properties in elements based on atomic number.
Key Properties:
- Atomic sizes, ionization energies, electron affinities.
Effective Nuclear Charge (Z_eff):
- Given by the formula: $Z_eff = Z - S$ where Z is atomic number and S is shielding effect by inner electrons.

Atomic Radius: Male decreases across a period (due to increased Z_eff) and increases down a group.
Ionization Energy: Energy required to remove an electron, generally increases across a period and decreases down a group.
Electron Affinity: Energy change when an electron is added to a gaseous atom; typically exothermic and increases across a period with noted exceptions for Group 2A, 5A, and 8A.