![Core Notation
• Writing out the entire electron configuration can get
very long
• Standard notation can be shortened to core
notation by using the noble gas configuration
before the element as the core
• Ex. Write the standard and core notation for P
• Standard - P: 1s2 2s2 2p6 3s2 3p3
• Core - P: [Ne] 3s2 3p3](https://image.slidesharecdn.com/atomictheoryextension-141212025931-conversion-gate02/85/Atomic-theory-extension-28-320.jpg)
Atomic theory extension
- 2. More history… • Niels Bohr - Developed a model based on observation that energized electrons appear to “jump” from one energy level to the next, and release a discrete amount of energy (quanta) when returning to lower levels.
- 3. • Werner Heisenberg - Uncertainty principle: The more precisely the position of an electron is determined, the less precisely it’s momentum is known and vice versa.
- 5. • Erwin Schrödinger – Contributed his famous wave equation used to predict the shapes of atomic orbitals. http://youtu.be/uWMTOrux0LM
- 7. A New Model of the Atom The Quantum Mechanical Model • Quantum mechanics provides a mathematical description of much of the dual particle-like and wave-like behavior and interactions of energy and matter • In the 1920’s Heisenberg and Schrodinger used quantum mechanics to describe the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.
- 8. Atomic Orbitals and Energy Levels • Atomic orbitals represent location around the nucleus that an electron has the greatest probability of existing Electron density distribution for H • Orbital arrangement follows a predictable pattern on the periodic table grouped by energy levels.
- 9. Atomic Orbitals and Energy Levels • Energy level number (n) corresponds the the distance from the nucleus • n=1, n=2, n=3 … (n=1 is closest to nucleus and lowest energy) • Within the energy levels there are sublevels containing orbitals • Each orbital can hold up to two electrons
- 10. Atomic Orbitals and Energy Levels • Definitions: • Shell – The set of orbitals having the same n-value • Ex. The third shell contains 3s, 3p & 3d orbitals • Sublevel (subshell) – The set of orbitals of the same type • Ex. The five 3d orbitals in the third shell are a sublevel
- 11. Types of Atomic Orbitals • Lowest energy orbitals are the s orbitals • Spherical shape around the nucleus • Increase in size as n increases • One s orbital found at each sublevel (2e-)
- 12. • Areas between orbitals are called nodes. • Nodes are regions of space where the probability of finding electrons is zero
- 13. Types of Atomic Orbitals • Second lowest energy orbitals are the p orbitals • Aligned along perpendicular axes • Increase in size as n increases • Three p orbitals found at each sublevel (6e-)
- 14. Types of Atomic Orbitals • Third lowest energy orbitals are the d orbitals • Complicated orbital shape • Increase in size as n increases • Five d orbitals found at each sublevel (10e-)
- 15. Types of Atomic Orbitals • Highest energy orbitals are the f orbitals • Complicated orbital shape • Increase in size as n increases • Seven f orbitals found at each sublevel (14e-)
- 16. Electron Configuration Rows in the periodic table correspond to the different energy levels and elements are grouped based on the type of orbitals their valence electrons are stored in
- 17. Electron Configuration • Three guiding rules: 1. Aufbau Principle “build up” 2. Hunds Rule “empty bus seat rule” 3. Pauli Exclusion Principle “opposite spins” • Remember: • Two electrons per orbital • 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals • For neutral atoms: Atomic number = protons = electrons
- 18. Electron Configuration 1. Aufbau Principle • Build Up • When filling atomic orbitals with electrons, the lowest energy orbitals are filled first.
- 19. Electron Configuration Standard Notation • Ex. Write the electron configuration of Nitrogen • Nitrogen has 7 electrons: 2 electrons in the 1s orbital, 2 electrons in the 2s orbital, 3 electrons in the 2p orbitals N 1s2 2s2 2p3
- 20. Electron Configuration 2. Hund’s Rule • Empty bus seats fill first • For same energy sublevels one electron is placed in each orbital before doubling up • Ex. Electron configuration for Carbon
- 21. Electron Configuration 3. Pauli Exclusion Principle • No two electrons can occupy the same space • Partnered electrons must have opposite spin • Possible values for spin are indicated with up or down arrows • Ex.
- 22. Tools:
- 23. Tools:
- 24. Practice • Write the electron configuration in standard notation for the following atoms: 1. Helium 2. Beryllium 3. Oxygen 4. Sodium 5. Sulphur 6. Calcium
- 25. Practice • Write the electron configuration in standard notation for the following atoms: 1. Helium: 1s2 2. Beryllium: 1s2 2s2 3. Oxygen: 1s2 2s2 2p4 4. Sodium: 1s2 2s2 2p6 3s1 5. Sulphur: 1s2 2s2 2p6 3s2 3p5 6. Calcium: 1s2 2s2 2p6 3s2 3p6 4s2
- 26. Practice • Show the electron configuration and spin states for the following atoms: 1. Lithium 1. Boron 1. Fluorine
- 27. Practice • Show the electron configuration and spin states for the following atoms: 1. Lithium: 1s2 2s1 ____ ____ 1. Boron: 1s2 2s2 2p1 ____ ____ ____ ____ ____ 1. Fluorine: 1s2 2s2 2p5 ____ ____ ____ ____ ____
- 28. Core Notation • Writing out the entire electron configuration can get very long • Standard notation can be shortened to core notation by using the noble gas configuration before the element as the core • Ex. Write the standard and core notation for P • Standard - P: 1s2 2s2 2p6 3s2 3p3 • Core - P: [Ne] 3s2 3p3