1.01.4 Electronic Structure Btec

Objectives

  • Objective 1: Explain the electronic configuration of Cr and Cu

  • Objective 2: Describe electronic structure in terms of S, P, and D orbitals

Sharp Recall Starter

  1. Draw the electronic structure of Na, S, & Ne

  2. State the 4 stages of mass spectroscopy

  3. Describe electron impact ionization

  4. Define Mr

Page 1 - Basic Concepts

  • Review of electronic structures, mass spectroscopy, and electron ionization

  • Importance of electron configuration in transition metals like Cr and Cu

Page 2 - Subatomic Particles

  • State the relative mass and charge of subatomic particles:

    • Protons: +1 charge, 1 amu

    • Neutrons: 0 charge, 1 amu

    • Electrons: -1 charge, 0.0005 amu (negligible)

  • Calculate proton, neutron, and electron counts in:

    • Na (11 protons, 12 neutrons, 11 electrons)

    • Cu (29 protons, 34 neutrons, 29 electrons)

    • P (15 protons, 16 neutrons, 15 electrons)

    • O2- (8 protons, 8 neutrons, 10 electrons)

    • Mg2+ (12 protons, 12 neutrons, 10 electrons)

  • Definition of an ion: An atom or molecule with a net electric charge due to the loss or gain of one or more electrons.

Page 3 - Electronic Structure Overview

  • Transition from GCSE to A-Level understanding of Electronic Structure

    • S, P, and D blocks require deeper knowledge

  • Writing Electron Configuration:

    • Example for Cr and Cu will demonstrate exceptions in configurations

  • Relation to Paper 1 Specification 3.1.1.3 reflecting fundamental understanding

Page 4 & 5 - Bohr Model and Energy Levels

  • Bohr model states electrons exist in defined orbits around the nucleus

    • Example: Calcium (Ca) configuration: 2.8.8.2.

  • Brief mention that A-Level covers more sophisticated models

Page 6 - Electron Shells and Capacity

  • Definition of energy levels (electron shells) by principal quantum numbers:

    • n = 1, 2, 3...

  • Maximum electrons in shells given by formula: 2n²

    • Shell 1 holds 2, shell 2 holds 8, shell 3 holds up to 18, and shell 4 holds 32 electrons.

Page 7 - Orbitals Structure

  • Each shell divides into sub-shells:

    • S, P, D orbitals

  • Orbitals where electrons are most likely found (95% probability)

Page 8 - Orbital Types

  • S orbitals: spherical shape with a maximum of 2 electrons

  • P orbitals: dumbbell-shaped with a maximum of 6 electrons

  • D orbitals: more complex shapes, holding up to 10 electrons

Page 9 - Electron Capacity in Orbitals

  • Each orbital can accommodate 2 electrons

    • Orbital shapes are fundamental to electron configuration patterns

Pages 10-11 - Energy Level Breakdown

  • Energy levels detailed:

    • 1st level: only 1s orbital

    • 2nd level: includes 2s and 2p

    • 3rd level: includes 3s, 3p, and 3d

    • 4th level: includes 4s, 4p, etc.

Page 12 - Sub-level Properties

  • Overview of orbitals:

    • s: 1 orbital, 2 electrons

    • p: 3 orbitals, 6 electrons

    • d: 5 orbitals, 10 electrons

    • f: 7 orbitals, holding up to 14 electrons

Pages 13-14 - Energy of Orbitals

  • Higher the principal quantum number, higher energy orbitals generally

  • Notably, the filling order for Cr and Cu is affected due to energy levels

  • 4s orbital has lower energy than 3d due to overlapping effects

Page 15 - Correct Terminology

  • Important terminology related to shells, energies, subshells, orbitals, and electron spins noted

Page 16 - Writing Electron Configuration

  • Writing notation:

    • Configuration: 1s2

    • Structure: First number = shell; Letter = subshell; Superscript = number of electrons

Pages 17-19 - Examples and Notation

  • Example for Cl:

    • Cl: 17 e- is represented as 1s2 2s2 2p6 3s2 3p5

    • Orbital notation with arrows indicating electron spins e.g., filled boxes represent orbitals holding electrons

Pages 20-22 - Filling Rules

  • Aufbau Principle: Electrons fill lower energy orbitals first

  • Pauli Exclusion Principle: Each orbital holds a maximum of 2 electrons with opposite spins

  • Hund's Rule: Every orbital is singly occupied before pairing begins

Pages 23-25 - Noting Exceptions

  • Notable Electron Configurations:

    • Chromium: 1s2 2s2 2p6 3s2 3p6 4s1 3d5 (NOT 4s2 3d4)

    • Copper: 1s2 2s2 2p6 3s2 3p6 4s1 3d10 (NOT 4s2 3d9)

Pages 26-29 - Ions and Shorthand Configurations

  • Ion formation involves losing highest energy electrons (4s before 3d)

  • Shorthand configurations utilizing noble gas notation, e.g., [Ar] 4s2.

Pages 30-34 - Block Placement

  • Elements classified into s, p, and d blocks based on their outer electron orbital

  • Importance of recognizing these blocks for understanding periodic trends.