Ch 4

1. Ch 4 Problem Set: pg 124-125 8,16 all, 18a-e, 19ab, 21ab, 26a, 27a, 28a, 30-33 all, 35 all, 38 all

2. 4.1 Refinements of the atomic model • Models of the atom so far: • Dalton – atoms are like little “bb’s” or solid spheres ( - then the electron gets discovered • Thomson – atom is like a charged “bb” (plum pudding model) • Rutherford - Gold foil experiment – hollow charged “bb” (atom is mostly empty space) • 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 • View vision learning example of hydrogen and helium atoms. • http://web.visionlearning.com/custom/chemistry/animations/CHE1.2-an-atoms.shtml

3. 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 (f). • C = λf • Wavelength (definition) = the distance between peaks of a wave • Light through prism leads to high energy (violet) low energy (red)

4. The Bohr Atom (1913) In 1913, Neils Bohr, a Danish physicist proposed: • 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, the less energy it has. • The farther the electron is from the nucleus, the more energy it has.

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

6. The Electromagnetic Spectrum • ROYGBIV - colors of the visible spectrum • bright line spectrum (bls) - frequencies of light give off by certain substances when energy is added to them. • heat sodium - yellow light 2 separate colors both in the yellow range • heat lithium - red light 6 colors (most prominent is red so overall appearance is red. • 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. 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 excited .

8. Bohr’s Proposal • 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 • Ground State Analogy on next slide

10. Ground State Analogy Filling Orbitals: the Apartment House Analogy To figure out where all of the electrons in a ground-state atom are located we can follow the set of rules given in the table below. Next to these rules are some rules that are intended to be similar in spirit. The idea is that filling up the orbitals in the shells of an atom is like renting out apartments in a large building. That is, provided that the superintendent of the building has to follow a very strange set of rules.

11. Bohr’s Energy Absorption Process:

12. Bohr’s Energy Absorption • 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):

13. Atomic Line Spectra • 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

14. The Bohr Model - Summary • When an atom absorbs energy, its electrons are promoted to a higher energy level. When the electron drops back down, energy is given off in the form of light. • Each distance fallen back is a specific energy, and therefore, a specific color. • Since electrons can fall from level 5 to 4, 5 to 3, etc., many colors are produced. • Click for animated H-atom: • Bohr's Atom: Quantum Behavior in Hydrogen - http://www.visionlearning.com/library/module_viewer.php?mid=51&l=15

15. The Bohr Model - Summary

16. More Bohr • 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.

17. Electron Configurations for Selected Elements Lithium Oxygen Fluorine Sodium The number of e- per shell = 2n2 (where n is then shell number) Animated Example, Atomic structure animation table - http://web.visionlearning.com/custom/chemistry/animations/CHE1.3-an-atoms.shtml

18. Electron Shells

19. Draw Bohr Models for the elements with atomic numbers 1-10 below then abbreviate with nucleus and numbers on rings

20. Bohr Model illustrations for elements 1-20 on the periodic table examples: 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.

21. Valence Shell Electron Patterns • 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

22. In sum, Bohr made 2 huge contributions to the development of modern atom theory (IMPORTANT) • 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 Ch4 worksheet #1 – question #1

23. 4.2 Quantum numbers and atomic orbitals & 4.3 Electron Configuration • The Bohr model describes the atom as having definite orbitals occupied by electron particles. • As with all chemistry, we soon learn that the Bohr model is a lie.

24. Schrödinger (1926) introduced wave mechanics to describe electrons • 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. • So, the new model really is wave mechanical model

25. There are really many types of orbitals – we can see them on the periodic table

26. Draw spdf blocks on blank periodic table Type(sublevel) number of orbitals total # of electrons shape S S 1 2 sphere P P 3 6 peanut D D 5 10 dumbbells F 7 14 flower

27. Quantum numbers show the “addresses” of electrons – each electron has 4 different quantum numbers: • principle (n): what shell, level, the e- is in n = 1,2,3...7 Principal Quantum Number • 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)

28. Orbital notation is another way to represent electron arrangement in atoms. • Electrons enter orbitals in a set pattern. For the most part, they follow these rules (3): 1. The Aufbau Principle - electrons must fill lower energy levels before entering higher levels.

29. More Orbital Notation 2. Hund’s Rule (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

30. Orbital Notation • Orbitals are like "rooms" within which electrons "reside". • The s subshell (or sublevel) has one s-orbital. The p subshell (sublevel) has three p-orbitals. • Each orbital can hold at most 2 electrons • See a good online illustration at http://www.avogadro.co.uk/light/aufbau/aufbau.htm

31. Orbital Notation/Electron Configuration • We can also do shorthand orbital notation (Noble gas plus the outer shell only) • Ca N Fe • Electron Configuration - a representation of the arrangement of electrons in an atom – is really just a variation of orbital notation. For H, instead of 1s1 • For He, instead of 1s2 • For C, instead of 1s2 2s2 2p2

32. EC for higher atomic #’s • At atomic # 19 (z = 19), a break in the pattern ensues. One would expect that the orbital to fill after 3p would be 3d, but alas, it is not. 4s is the next level we fill as it has lower energy than 3d. The same is true for calcium. (Similar to the Bohr diagrams.) • Now the next level to fill is 3d, but we have exception columns. • Cr is 4s1 3d5 and Cu is 4s13d10 - students should be able to identify these elements simply based on how many electrons they have.

33. Shorthand Electron Configuration • We can also do shorthand electron configuration. Use previous row’s noble gas in brackets, then present the electron configuration for the current row. • Be Na Si

34. Significance of electron configurations • Valence shell electrons – electrons in the highest occupied energy level • no atom has more than 8 • Noble gases - 8 valence electrons – least reactive of all elements Noble Gas Notation • Lewis Dot structures: NNESWESW (cheating) also show correct way, count to 8

35. Lewis Dot Diagrams (EDD) Exception – He Try Lewis Dot structures for additional elements here: