Atomic structure

Dalton’s Atomic Theory

• All Matter is Atoms: All matter consists of very small particles called atoms.

• Identical Atoms: Atoms of a given element are identical in properties; atoms of different elements are different.

• Indivisibility: Atoms cannot be subdivided, created, or destroyed. Chemical Compounds: Atoms combine in simple whole-number ratios to form compounds.

• Rearrangement in Reactions: In chemical reactions, atoms are combined, separated, or rearranged.

Modern Atomic Theory

• Divisible Atoms: Atoms are not indivisible; they can be split into smaller particles.

• Isotopes Exist: Elements can have atoms with varying masses (isotopes).

• Fundamental Concepts: Key points from Dalton's theory remain valid:

  • All matter is made of atoms.

  • Atoms of different elements differ in properties.

Discovery of the Electron

• Cathode Rays: Experiments revealed a negatively charged particle, the electron, within atoms.

• Negative Charge: Cathode rays bend towards positive charges, indicating they are negatively charged.

Plum Pudding Model

• Proposed by J.J. Thomson: Electrons are distributed within positively charged material, like raisins in a cake.

Gold Foil Experiment

• Conducted by Ernest Rutherford:

  • Positively charged alpha particles were directed at gold foil.

  • Hypothesis: If the plum pudding model were correct, most particles would pass through without deflection.

Atomic Structure Insights

• Proton Count: Different elements have unique proton numbers.

• Atomic Number (Z): Represents the number of protons in an atom.

• Isotopes: Variants of the same element with different neutron counts but the same number of protons and electrons.

Atomic Number and Mass Number

• Mass Number (A): Total of protons and neutrons in the nucleus.

• Example: A helium nucleus with 2 protons and 2 neutrons has an atomic number of 2 and a mass number of 4.

Isotopes and Properties

• Isotopes: Same protons/electrons, varying neutrons.

• Most elements are mixtures of isotopes.

Relative Atomic Masses

• Atomic Mass Unit (amu): Defined as 1/12 the mass of a carbon-12 atom.

• Average Atomic Mass: Weighted average based on natural isotope composition.

• Example Calculation for Copper:

  • Copper-63: 69.15% with mass 62.929601 amu

  • Copper-65: 30.85% with mass 64.927794 amu

  • Calculation: $(0.6915 imes 62.929601 + 0.3085 imes 64.927794) = 63.55$ amu

Bohr Model of the Hydrogen Atom

• Proposed by Niels Bohr:

  • Electrons orbit the nucleus at fixed distances (orbits).

  • Photon Emission: Electrons emit photons when transitioning to lower energy levels.

  • Absorption: Electrons gain energy to jump to higher orbits.

Schrödinger Wave Equation

• Developed by Erwin Schrödinger in 1926:

  • Electrons are described as waves, leading to modern quantum theory.

  • Orbitals: Electrons occupy probabilistic regions rather than fixed orbits.

Quantum Numbers and Atomic Orbitals

• Principal Quantum Number (n): Indicates the main energy level.

• Angular Momentum Quantum Number (l): Defines orbital shape.

• Magnetic Quantum Number (m): Determines orbital orientation.

• Spin Quantum Number: Indicates electron spin state (+1/2 or -1/2).

Electron Configuration Rules

• Aufbau Principle: Electrons fill lower-energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund’s Rule: Electrons occupy all degenerate orbitals singly before pairing up.

Orbital Notation

• Unoccupied orbitals are shown with lines; electrons represented with arrows.

• Example Configurations:

  • Unoccupied: |

  • One electron: ↑

  • Two electrons: ↑↓

Noble-Gas Notation

• Refers to electron configurations for Group 18 elements (noble gases).

• Atoms often aim to achieve stable configurations resembling noble gases, usually with eight electrons in the outer shell.