Chapter 1: Atomic Structure

Chapter 1: Atomic Structure

Subatomic Particles

- Protons carry a positive charge (+1), have a mass of approximately 1 amu, and reside in the nucleus.

- Neutrons are neutral, also about 1 amu, and found in the nucleus.

- Electrons have a negative charge (–1), a much smaller mass (1/1836 amu), and orbit the nucleus in shells.

Isotopes

- Atoms of the same element with identical proton numbers but different neutron counts.

- Example: Carbon-12 vs. Carbon-14.

- They behave the same chemically but differ in physical properties like density and mass.

Atomic and Mass Numbers

- Atomic number (Z) = number of protons.

- Mass number (A) = number of protons + neutrons.

Relative Masses

- Relative atomic mass (Ar) is the weighted average of isotopes compared to 1/12 of carbon-12.

- Relative isotopic mass refers to a specific isotope’s mass on the same scale.

Mass Spectrometry

- Used to determine isotopic composition and calculate Ar.

- Key steps: Ionization → Acceleration → Deflection → Detection.

- Produces a mass spectrum showing abundance vs. mass/charge ratio.

Electronic Configuration

- Electrons fill orbitals in order of increasing energy.

- Aufbau Principle: Fill lowest energy orbitals first.

- Pauli Exclusion Principle: Max two electrons per orbital, opposite spins.

- Hund’s Rule: Orbitals of equal energy get one electron each before pairing.

- Example: Oxygen (Z=8) → 1s² 2s² 2p⁴

Ion Formation

- Cations form by losing electrons → positive charge.

- Anions form by gaining electrons → negative charge.

- Electron configurations adjust accordingly.