Atomic Structure Flashcards
Introduction to Atomic Structure
Atomic structure refers to the arrangement of subatomic particles—protons, neutrons, and electrons—within an atom. Understanding the atomic model is fundamental for explaining chemical properties and behaviors of elements and compounds.
• Atoms: The smallest units of matter that retain the properties of an element.
• Subatomic particles:
• Protons: Positively charged particles found in the nucleus.
• Neutrons: Neutral particles found in the nucleus.
• Electrons: Negatively charged particles found in energy levels (shells) surrounding the nucleus.
The Atomic Model
2.1 Early Atomic Models
• John Dalton (1803): Proposed that atoms are indivisible and indestructible spheres.
• J.J. Thomson (1897): Proposed the “Plum Pudding Model,” where electrons were embedded in a positively charged “pudding.”
• Ernest Rutherford (1911): Introduced the “Nuclear Model,” proposing that most of the atom’s mass and positive charge are concentrated in a tiny nucleus, with electrons orbiting the nucleus.
• Niels Bohr (1913): Suggested that electrons occupy fixed energy levels around the nucleus, with specific orbits.
2.2 Modern Atomic Theory (Quantum Model)
• Electron Cloud Model: The current model of the atom describes electrons as existing in orbitals, which are regions of space with a high probability of finding an electron. These orbitals correspond to specific energy levels.
The Structure of the Atom
3.1 Subatomic Particles and Their Properties
• Protons- Mass: 1 (atomic mass unit) Charge: +1
• Neutrons - Mass: 1 and Charge: Neutral
• Electrons- Mass: approx 1/1836 Charge: -1
3.2 Atomic Number and Mass Number
• Atomic Number (Z): The number of protons in an atom, which defines the element.
• Mass Number (A): The sum of protons and neutrons in the nucleus.
• Isotopes: Atoms of the same element with different numbers of neutrons, resulting in different mass numbers.
Electron Configuration
Electron configuration describes the arrangement of electrons in an atom’s energy levels or orbitals. Electrons occupy the lowest energy orbitals first, as described by the Aufbau principle.
4.1 Principles Governing Electron Configuration
• Aufbau Principle: Electrons fill orbitals starting from the lowest energy level to the highest.
• Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons, and these must have opposite spins.
• Hund’s Rule: Electrons fill degenerate orbitals (orbitals of the same energy) singly before pairing up.
4.2 Electron Shells and Subshells
• Electron Shells: The main energy levels (n = 1, 2, 3, 4, etc.).
• Subshells: Within each shell, there are subshells (s, p, d, f), which hold a specific number of electrons:
• s: 2 electrons
• p: 6 electrons
• d: 10 electrons
• f: 14 electrons
4.3 Example: Electron Configuration of Oxygen
• Oxygen (Z = 8): The electron configuration is 1s^2 2s^2 2p^4.
4.4 Transition Metals and d-Block Elements
• In transition metals, electrons fill the 3d orbitals after the 4s orbitals, which results in exceptions to the expected order of filling.
Ionisation Energy
Definition of Ionisation Energy
• Ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.
5.2 Factors Affecting Ionisation Energy
• Atomic Radius: A larger atomic radius means the outer electrons are farther from the nucleus and experience weaker attraction, making ionisation easier.
• Nuclear Charge: More protons in the nucleus result in a stronger positive charge, making it harder to remove electrons.
• Electron Shielding: Electrons in inner shells can shield outer electrons from the full nuclear charge, decreasing the ionisation energy.
5.3 Trends in Ionisation Energy
• Across a Period: Ionisation energy increases across a period due to an increase in nuclear charge and a decrease in atomic radius.
• Down a Group: Ionisation energy decreases down a group because of increased electron shielding and a larger atomic radius.
Mass Spectrometry
Mass spectrometry is a technique used to determine the relative isotopic mass and abundance of isotopes in a sample.
6.1 How Mass Spectrometry Works
• Ionization: Atoms are ionized by knocking off electrons.
• Acceleration: The ions are accelerated by an electric field.
• Deflection: The ions are deflected by a magnetic field based on their mass-to-charge ratio (m/z).
• Detection: The ions are detected, and their abundance is recorded.
6.2 Key Concepts in Mass Spectrometry
• The mass spectrum produced shows peaks corresponding to the different isotopes of an element and their relative abundances.
• Relative Atomic Mass (Ar) is calculated by:
{Ar} = {sum of (isotopic mass x abundance)} {total abundance}}
Exam Tips:
Ensure you can draw and explain electron configurations for various elements, especially transition metals.
• Be familiar with periodic trends, as they form a significant part of exam questions.
• Practice using mass spectra to calculate relative atomic masses.
