1. Electron Configuration And a Brief Introduction to the Quantum Model of the Atom

2. Bohr’s Atomic Model  Electrons are located in specific energy levels n=1 n=2

3. Energy Levels and the Periodic Table  The main energy level number corresponds to a period in the periodic table of the same number  All the elements of that period use that energy level for their electrons n=1 Period 1 n=2 Period 2 n=3 Period 3 etc.

4. Sublevels  Sublevels are located within an energy level  Each sublevel has a name Energy LevelNames of Sublevels 1 st Energy Level; n=1s 2 nd Energy Level; n=2s and p 3 rd Energy Level; n=3s, p, and d 4 th Energy Level; n=4s, p, d, and f

5. Subdivisions in the Periodic Table  The periodic table can be subdivided to show the sublevels

6. Orbitals  Energy levels can have different sublevels  Sublevels can have different orbitals  Orbitals are located inside sublevels  Different sublevels have different numbers of orbitals Sublevel# of Orbitals Possible s1 p3 d5 f7

7. Orbitals  Only 2 electrons can fit in each orbital  There are 2 electrons in an s orbital  There are 2 electrons in a d orbital  Since there are 5 d orbitals, there are a total of 10 electrons in the d sublevel  There is a slight variation in the energy of electrons between sublevels, but electrons in orbitals of the same sublevel have the same energy https://www.wisc-online.com/learn/natural-science/chemistry/gch904/the-structure-of-an-atom

8. Diagrams of s, p, and d

9. Electron Configuration  Each element has a distinct electron configuration that can be written using the energy levels, sublevels, and orbitals that its electrons occupy  The electron configuration for hydrogen is 1s11s1 Main energy level sublevel Indication Of Orbital

10. Electron Configuration  Unfortunately, some orbitals at higher energy levels fill before all orbitals at lower energy levels  This makes an atom more stable  For example, 4s fills before 3d

11. Using the Periodic Table  We can use and element’s position on the periodic table to determine its electron configuration  The electron configuration for oxygen is: 1s 2 2s 2 2p 4

12. Exception  A d sublevel that is half full or full (i.e. 5 or 10 electrons) is more stable than the s sublevel of the next energy level  As a result, electron configurations rarely end in __ d 4 or __ d 9  An electron is taken from the previous s sublevel to change this to __ d 5 or __ d 10

13. Example: chromium The electron configuration for chromium is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 not 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4

14. Abbreviated Electron Configuration  The abbreviated electron configuration only shows the number of electrons in each main energy level  The abbreviated electron configuration for oxygen is: 2, 6  The abbreviated electron configuration for chromium is: 2, 8, 13, 1

15. The Quantum Model of the Atom Bohr’s Work Bright Line Spectra Electron Transitions

16. Properties of Light Light as a Wave Light can be thought of as electromagnetic radiation having a particular wavelength and frequency Light as a Particle Albert Einstein proposed almost a century ago that electromagnetic radiation can be viewed as a stream of particles known as photons A photon has a particular amount of energy associated with it

17. Back to Bohr Neils Bohr wanted to determine exactly where electrons were located in an atom Studied gaseous hydrogen atoms and the specific colours of light they produced when electricity was applied to them When Bohr focused the light through a prism, he observed lines of only certain specific colours These lines are known as a bright line spectrum Every element on the periodic table has its own unique bright line spectrum

18. Bright Line Spectra Bohr concluded that electrons exist in specific energy levels in the atom and that these energy levels are quantized (i.e. have a certain value of energy associated with them the energy levels where electrons are normally found are called ground states If an electron absorbs sufficient energy it moves to a higher energy level to produce an excited state When the electron releases the energy, it drops back to a lower energy level, and the energy is released in the form of electromagnetic radiation

19. The Electromagnetic Spectrum

20. The wavelength of the emitted light indicates the difference in energy between the ground state and the excited state Each wavelength corresponds to a specific type of electromagnetic radiation, which may or may not be visible Electron Transitions

21. Electron Transitions in a Hydrogen Atom INVISIBLE SEEN AS 4 DIFFERENT COLOURS

22. The Visible Light Spectrum

23. The Bright Line Spectrum for Hydrogen While many electron transitions are possible for a hydrogen atom, only four of them produce visible light

24. Inside the Hydrogen Atom 1. Electron jumps from n=1 to n=4 2. Electron jumps from n=4 to n=2. 3. Light is emitted (486nm)

25. Who Cares? (besides physicists and chemists) The absorption and emission of electromagnetic radiation is one of the most powerful tools used to probe molecular structure and chemical reactions It forms the basis of nuclear magnetic resonance imaging (NMR) It is intrinsic to many analytical techniques used to monitor manufacturing and the environment Trace materials (evidence from a crime scene, lead in paint, mercury in drinking water) can be identified