1. Through the work of Einstein, Planck, DeBroglie, Bohr, Schrodinger, Heisenberg we know Depending on the experiment, light has particle like behavior (an example…photoelectric effect) Particles (the electron) has wave like behavior (example: the emission spectrum) Our goals in the first part of this chapter is 1.To describe light and particles in terms of energy and wavelength, apply DeBroglie’s relationship and the photoelectric effect 2.To describe likely positions of where the electron is located to understand chemical behavior.

2. Light is a form of electromagnetic radiation Electromagnetic radiation is energy (radiant) that as it travels consists of electric and magnetic fields Travel in waves that can be described by their wavelength (lamda –  and frequency ( - nu) In a vacuum, travel at the speed of light, c = 3 x 10 8 m/s = c

3. White light is made up of light waves of different wavelengths Disperses light into spectrum of colors. There are many lines that blend together producing a continuous spectrum ( 400-700 nm).

4. Hydrogen

5. Neils Bohr Proposed electrons travel in successively larger orbits and when an electron jumps from an outer orbit to an inner one, it emits light. Proposed the planetary model - suggested the electron orbits the nucleus as planets orbit the sun orbitals Bohr theorized that the lines in the hydrogen spectrum corresponded to certain quantum of energy emitted or absorbed when the electron moved from one “orbit” to another.

6. Spectroscopy Method of studying substance that are exposed to exciting energy. Can be used to identify substances, since each element emits a unique collection of lines. The presence of spectral lines of specific frequencies suggests that the energy of the electron in an atom is restricted to a series of discrete values

7. Max Planck Proposed energy is quantized Energy can only be lost or gained in packets Each packet is called a quantum Energy of quantum of electromagnetic radiation is proportional to frequency,  E = h 

8. Indium compounds give a blue- violet flame test. The atomic emission responsible for this blue-violet color has a wavelength of 451 nm. Calculate the energy of a single photon of this wavelength.

9. Albert Einstein Used Planck’s equation to explain photoelectric effect Light consists of quanta of energy that behave like tiny particles of light(photon). Electrons can absorb energy from photons, but they follow an "all or nothing" principle. All of the energy from one photon must be absorbed and used to liberate one electron. 1/2mv 2 = h -  1/2mv 2 : Kinetic energy of electron ejected h : Energy of incident radiation  : Work function, minimum energy needed to eject electron

10. The photoelectric work function for Mg is 5.90 x 10 -19 J. Calculate the minimum frequency of light required to eject electrons from Mg. Light of wavelength 285 nm shines on a piece of Mg metal. What is the speed of the ejected electron?

11. Albert Einstein Used Planck’s equation to explain photoelectric effect In related development, proposed special theory of relativity E = mc 2 Light consists of quanta of energy that behave like tiny particles of light(photon). Electrons can absorb energy from photons, but they follow an "all or nothing" principle. All of the energy from one photon must be absorbed and used to liberate one electron.

12. What were the contributions of DeBroglie? Heisenberg? Schrodinger? Matter Waves, Uncertainty Principle – location and momentum of particle are complimentary; can’t both be known simultaneously with precision; can’t specify precise location of particle if it behaves like a wave Developed an equation that describes the wavelike properties of matter, we use the wave function to express the probability of finding the electron in a specific region in space (orbital), Quantum number used to describe the orbital and the electron = h / mv

13. 1.Probability of finding electron at different points is calculated Some points will have higher probability than others 2.If connect all points of high probability, three dimensional shapes are formed 3.The most probable place to find the electron will be some place in that shape

14. Quantum Numbers Set of four numbers that completely specify each electron 1. Prinicipal quantum number, n, corresponds to the energy level 2. Second quantum number, l, corresponds to the sublevel within the energy level. It provides information about the shape of the electron cloud 3. Third quantum number, m, is the orbital quantum number and describes the orientation in space of orbital 4. Fourth quantum number, s, is the spin quantum number and describes rotation of an electron on its axis. Values are ½ or - ½

15. SublevelOrientations (Orbitals) s p d f Region around nucleus where electrons are likely to be found. Each orbital holds two electrons that have opposite spin

16. 1s __ 2s __ 2p __ __ __ 3s __ 3p __ __ __ 4s __ 3d __ __ __ __ __ 4p __ __ __ 5s __ 4d __ __ __ __ __ 5p __ __ __ 6s __ 4f __ __ __ __ __ __ __ 5d __ __ __ __ __ 6p __ __ __ 7s __ 5f __ __ __ __ __ __ __ 6d __ __ __ __ __

17. Atomic Orbitals Region around nucleus where electrons are likely to be found Energy LevelSublevels 1s 2s, p 3s, p, d 4s, p, d, f Each orbital holds two electrons that have opposite spin

18. Follow 3 rules to configure the electrons 1. Aufbau Principle - electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states 2. Pauli Exclusion Principle - two electrons cannot share the same set of quantum numbers within the same system. Therefore, there is room for only two electrons in each orbital and the electrons have opposite spin. 3. Hund’s Rule – in equal energy orbitals, arrange the electrons to achieve the maximum number of unpaired electrons.

19. DO NOW Draw the orbital notation for arsenic. How many unpaired electrons does it have? Compare paramagnetic and diamagnetic. Write the electron configuration for manganese, Br -, Ba 2+ Depict an atom (Bohr Model) of Uranium-235. Paramagnetic has net unpaired spins and it attracted by a magnet. Diamagnetic does not.

20. Take out a periodic table and paper only. 1.Write out the electron configuration for osmium-191 2.Write out the electron configuration for A) Na and Li B) O and S C) He, Ne, Ar For each grouping in number 2, do you notice any similarities?

21. Write out the electron configuration for Lithium Oxygen Cesium Abbreviated Electron Configuration

22. What is a trend ? Atomic Radius Ionic Radius Ionization Energy Electron Affinity

23. Atomic Radius Distance from nucleus to outermost electron Down a Group- Increases Across a Period – Decreases

24. Positive ion smaller than atom Negative ion larger than atom Distance from nucleus to outermost electron in an ion Ionic Radius ATOM + - < << <

25. Ionization Energy Minimum energy required to lose an electron and form a positive ion Down a Group- Decreases Across a Period – Increases

26. Successive Ionization Energies Look at the chart on p. 310 If you scan Magnesium’s ionization energies a large jump is observed between the 2 nd and 3 rd. Explain. Where would you predict a large jump in successive ionization energies for Sr?

27. Practice! 1. Compare the sizes of… (Explain your trend) a.Chlorine and magnesium b.Nitrogen and bismuth 2. What is the charge of the ion formed by a. nitrogen? Also compare the radius of the nitrogen atom to the nitride ion.. b. magnesium? Compare magnesium ion and atom radii. 3. Compare ionization energy of____. Which is larger and why? a. strontium and antimony b. sodium and francium

28. Electron Affinity Energy to gain an electron and form a negative ion Down a Group- Increases Across a Period – Decreases

29. Electronegativity Ability to attract electrons in a chemical bond Down a Group- Decreases Across a Period – Increases