1. Chapter 4 Arrangement of Electrons in an Atom

2. 4.1 Refinements of the atomic model • Models of the atom so far: • Dalton – atoms are like little “bb’s” - then the electrongets discovered • Thomson – atom is like a charged “bb” • Rutherford - Gold foil experiment – hollow charged “bb” • Bohr model of the atom (1913) – Neils Bohr – Danish Physicist • The Bohr model of the atom comes from the idea that light is waves of energy http://web.visionlearning.com/custom/chemistry/animations/CHE1.2-an-atoms.shtml

3. The Bohr Atom (1913) • All the positive charge was in the nucleus • Electrons orbited the nucleus much like planets orbit the sun (at fixed distances) • The closer the electrons to the nucleus, theless energy it has. • The farther the electron is from the nucleus, the more energy it has.

4. The Electromagnetic Spectrum • Visible light, x-rays, ultraviolet radiation, infrared radiation, microwaves and radio waves are all part of the electromagnetic spectrum

5. The Electromagnetic Spectrum • The spectrum consists of electromagnetic radiation – energy that travels like a wave • Waves can be described by the wave equation which includes velocity (c = speed of light), wavelength (λ) and frequency (ν). • Wavelength (definition) = the distance between peaks of a wave • Light through prism leads to high energy (violet) low energy (red)

6. The Electromagnetic Spectrum • ROYGBIV - colors of the visible spectrum • Bright Line Spectrum (BLS) – caused by e- emitting energy as they return tolower energy levels energy level. • heat sodium - yellow light 2 c • heat lithium - red light • elements can appear to give off the same color light, but each will have its own BLS • BLS - used to determine identity of an element • BLS - validates Bohr’s idea that electrons jump to different energy levels and give off different wavelengths of light

7. The Electromagnetic Spectrum • Light from the sun (white light) appears as a continuous spectrum of light. • Continuous Spectrum of Light (definition) = There are no discrete, individual wavelengths of light but rather all wavelengths appear, one after the other in a continuous fashion • Spectroscopy (definition) = the study of substances from the light they emit. • We will use spectroscopes (An instrument that splits light into its component colors) and flame tests to study elements because each element emits a different spectrum of light when exited .

8. Birght Line Spectrum • Bohr proposed that the energy possessed by an e- in a H- atom and the radius of the orbit are quantized (bls) • Quantized (definition): a specific value (of energy) Like a set of stairs, the energy states of an electron is quantized – i.e. electrons are only found on a specific step The ramp is an example of a continuous situation in which any energy state is possible up the ramp

9. Bohr’s Energy Absorption Process • Light or energy excites an e- from a lower energy level (e- shell) to a higher energy level • These energy levels are “ quantized “ (the e- cannot be in between levels), the e- disappears from one shell and reappears in another • This absorption or excitation process is called a quantum leap or quantum jump

10. Bohr’s Energy Absorption Process • Ground State Analogy = a spring and two balls This is an energy emission process and what we observe in the hydrogen line spectrum Both the atom and e- now have higher energy The e- absorbs energy in the ground state and is excited to a higher level

11. Bohr’s Energy Absorption Process • When energy is added, the electron is found in the “excited state.” • The Excited State (definition) = an unstable, higher energy state of an atom • An illustration of Bohr’s Hydrogen atom (from ground to excited state):

12. Bohr’s Energy Absorption Process • The atomic line spectral lines - when an e- in an excited state decays back to the ground state The electron loses energy, light (colors) is emitted and the e- returns to the ground state This is another illustration of bls.

13. The Bohr Model - Summary

14. The Bohr Model - Summary • Bohr also predicted that since electrons would occupy specific energy levels and each level holds a specific number of electrons • The maximum capacity of the first (or innermost) electron shell is twoe-. • Any element with more than twoe-, the extra e- reside in additional electron shells.

15. The Bohr Model - Summary Electron Configurations for Selected Elements • The number of e- per shell = 2n2 (where n is the shell number)

16. Bohn Models Draw Bohr Diagrams for the elements 1-10. Save room to draw them short hand also

17. Short Hand Bohr Model • Write the symbol of the element • Use a ) to represent each shell • Write the # of e- in each shell • Ex.

18. The Truth About Bohr Models • At atomic # 19 (z = 19), there is a a break in the pattern. One would expect that energy level #3 would continue to fill up. However, the next two electrons go into the next energy level. Look at K and Ca:

19. The Truth Continued…… • So, there is a relationship between the main column # and the number of outershell electrons. • Column # = the number of valence electrons • And, there is a relationship between the row # and the number of energy levels. • Row # = the number of shells • The Bohr model truly works well for the H atom only • for elements larger than H the model does not work.

20. Bohr Summed Up • Bohr made 2 huge contributions to the development of modern atom theory • He explained the atomic line spectra in terms of electron energies • He introduced the idea of quantized electron energy levels in the atom • The Bohr atom lasted for about 13 years and was quickly replaced by the quantum mechanical model of the atom. The Bohr model is a good starting point for understanding the quantum mechanical model of the atom Do ws# 1, question 1- Use short-hand configuration

21. 4.2 Quantum Numbers and Atomic Orbitals • 4.3 Electron Configuration

22. Quantum Numbers & Atomic Orbitals • The Bohr model describes the atom as having definite orbitals occupied by electron particles. • Schrödinger (1926) introduced wave mechanics to describe electrons – proved Bohr’s Model to be a lie • Based his idea that electrons behaved like light (photons). • Electrons show diffraction (interference) properties like light. • Treats electrons as waves that are found in orbitals. • Orbitals (definition) = clouds that show region of probable location of a particular electron.

23. Wave Mechanical Model • The Bohr model really is the wave mechanical model • There are many types of orbitals – we can see them on the periodic table

25. Subatomic Orbitals S P D

26. Quantum Numbers • An electron’s address • principle (n): what shell, level, the e- is in n = 1,2,3...7 • azimuthal (l): energy sub level - s, p, d, f • magnetic – orientation of orbital about the nucleus (s has only 1, p has 3, etc.) • spin - clockwise or counterclockwise (+1/2 or -1/2)

27. Label Your Periodic Tabel • On your periodic table, shade azimuthals,p,d,f blocks different colors • Label the principal quantum numbers…1-7 • Label the valence electrons across the top

28. Electron Configuration • Electron Configuration - a representation of the arrangement of electrons in an atom

29. Electron Configuration • Examples of electron Configuration • 1. Li 1s22s1 • 2. C 1s22s22p6 # of e- in that shell principle azimuthal

30. Electron Configuration • Take note that after 4s is filled, 3d is than filled before 4p. • …… 6s than 4f than 5d than 6p • When writing out the electron configuration, always write your numbers in numerical order • Y 1s22s22p63s23p64s23d104p65s24d1 – NO! • Y 1s22s22p63s23p63d104s24p64d15s2

31. Electron Configuration • Examples: • Be • O • Ca • Mn

32. Electron Configuration • Examples • Pb • Os

33. Electron Configuration • Short Hand • Write the name of the last noble gas • Write the electron config. that follows • Ex. Fe [Ar]3d64s2 • Exceptions • Cr [Ar] 3d54s1 • Cu, Ag, Au- all s’s donate 1 e- to make the d orbital full • Cu [Ar] 4s13d10

34. Orbital Notation • Electrons enter orbitals in a set pattern. For the most part, they follow these rules: • 1) The Aufbau Principle - electrons must fill lower energy levels before entering higher levels.

35. Orbital Notation • Orbitals are like "rooms" within which electrons "reside". • The s subshell has one s-orbital. The p subshell has three p-orbitals. The d subshell has 5 and f has 7. • Each orbital can hold at most 2 electrons

36. Orbital Notation • 2. Hund’sRule (better known as the Bus Rule) • Before any second electron can be placed in a sub level, all the orbitals of that sub level must contain at least one electron – spread out the e- before pairing them up. • 3. Pauli Exclusion Principle - electrons occupying the same orbital must have opposite spin. • See a good online illustration at http://www.avogadro.co.uk/light/aufbau/aufbau.htm

37. Orbital Notation H 1s F 1s 2s 2p

38. Orbital Notation • Examples: • Li F • Na Sc

39. Orbital Notation • We can also do shorthand orbital notation (outer shell only) • Ca N • Fe Ag [Kr] 4d105s1 Ag [Kr] 4d 5s

40. Significance of Electron Configurations • Valence shell electrons - outermost electrons involved with bonding • no atom has more than 8 valence electrons • Noble gases - 8 valence electrons – least reactive of all elements • Lewis Dot structures: NSEW (cheating) also show correct way, count to 8