Unit 3, Part 2 : Electronic Configuration

1. Unit 3, Part 2: Electronic Configuration Light and the electron

2. 1 Radiant Energy summary • Light is a form of electromagnetic radiation. Its energy is transmitted in the form of waves. • The electromagnetic spectrum spans a wide range of wavelengths, from radio to X-rays.

3. Parts of a wave • Wavelength (λ)- the distance from crest to crest. • Frequency (f)- how many wavelengths per second. • Amplitude – the height of the wave.

4. Sunlight contains ultraviolet light. • X rays & gamma rays are used in medicine. • Infrared light can be felt as warmth. • Microwaves are used in cooking and communication. • Radio waves are used in communications and radar.

5. 2 Quantum Theory • The “classic” model (or wave model) of light was a great beginning to understanding light, however, it could not explain everything. • Ex. Why different ranges of light were emitted at different temperatures… why do different elements give off different colors when burned?

6. PLANCK’S THEORY • In 1900, Max Planck was able to predict how the emission spectrum of an object would change with temperature. He was able to make theses predictions after suggesting you could have a piece of radiation (light). “Planck proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs and he called each of these pieces of energy a quantum”. (a fixed amount) Planck’s theory creates a relationship between the frequency/wavelength and the energy of the object.

7. E=hv • E= amount of energy, v= the frequency and the h= 6.6262 x 10-34 j-s called Planck’s constant. • His theory suggests the energy absorbed or emitted by atoms is in discrete amounts (quanta), restricted to certain quantities.

8. THE PHOTOELECTRIC EFFECT • In 1905, Albert Einstein used Planck’s equation to explain another phenomenon- the Photoelectric Effect. This occurs when e- are ejected from the surface of metals when particular frequencies of light shine on them. Each metal requires a different frequency to release the e-s. Quantawere the answer to explain why this was occurring. Einstein proposed that the light consisted of quanta of energy and he named this type of energy photons.

9. How much energy does each photon carry?? Planck’s equation tells us. • How did Einstein explain this effect?? The photon transfer’s its energy to the electron in the metal. The electron must use only one photon and all the energy from that photon to “jump”. If there isn’t enough energy in the photon… the electron won’t move. Thus the photon with the perfect amount of energy will move the e- and no other! http://www.lewport.wnyric.org/mgagnon/Photoelectric_Effect/photoelectriceffect1.htm

10. DUAL NATURE OF RADIANT ENERGY • Einstein’s idea of the photon was proven by an American physicist, Arthur Compton. He showed that a photon colliding with an electron behaves the same way two marbles do when they collide. THUS… a photon behaves like a particle(a particle that travels at the speed of light with a specific frequency and wavelength). • LIGHT IS A WAVE-ICLE!!!!

11. 3 Another Look at the Atom • Remember… electrons can only gain and lose energy in chunks (quanta). • LINE SPECTRA: • A spectrum containing specific colors/wavelengths • Elements emit unique line spectra(like a fingerprint for elements). Can use these spectra to ID the elements in stars. • HOW??? Electrons absorbing photons are jumping and relaxing?

12. THE BOHR MODEL OF THE HYDROGEN ATOM • In 1911, the Danish physicist, Niels Bohr, was the first to make the connection between the line spectra and atomic structure. He took Rutherford’s idea of “orbitals of electrons around a nucleus” and applied Planck’s Quantum Theory to it and suggested:

13. *Spectral lines are generated by quantized energy of the electron(s) in the atom. • *Electrons in different “orbitals” have a different amounts of energy associated with them. • *The “orbitals” are assigned quantum #s, n, with the ground state = 1 & closest to the nucleus. When the electron absorbs the perfect amount of energy… it will jump to a higher orbital. The electron is now in the excited state (with larger orbits).

14. *The spectral line are generated when the electron relaxes back to it’s original orbital/energy state(the ground state).The energy it absorbed to jump is released and is = to the difference in energy between the two energy levels. • (ex from the pic. above… the difference between n=3 and n=2). • *The energy emitted is seen as a color of light according to Planck’s equation. • Bohr’s explanation was perfect for Hydrogen, and Hydrogen only! However, he established the important idea of quantized energy levels for electrons.

15. MATTER WAVES • If light can be a wave-icle… having properties of both waves and particles… then can MATTER behave in a similar fashion? Is matter a wave-icle too?

16. The Frenchmen, Louis de Brogliefirst suggested and proved it possible mathematically in 1924. And just like Planck it took American ingenuity to prove him correct. • In 1927, Clinton Davisson and Lester Germer, of Bell Labs, showed electrons had wavelike properties which could be predicted precisely by de Broglie’s equation. In fact… any moving object possesses wavelike behavior! Even you!

18. HEISENBERG’S UNCERTAINTY PRINCIPLE • Another important idea was presented in 1927 when Werner Heisenberg proposed his uncertainty principle. Simply stated: the position and momentum of a moving object cannot be measured and known at the moment. You can’t know where it is and it’s movement at the same time in order to make predictions. This is because of the wave-icle phenomenon. This has a big impact on small items like electrons; however, on large items like you and me, it does not. • How do you see the blade of a rotating fan- know its exact location- and know it’s momentum at the same moment? Electrons are far more difficult to assess. They move in more unpredictable patters and Heisenberg’s principle suggests it’s not even logical to think of electrons existing in “orbits” with distinct paths.

19. 4 A New Approach to the Atom • **Quantum-mechanical model** treating the electron as a wave of quantized energy, has regions of electron density and not orbitals giving electrons a distinct path/position.

20. PROBABILITY AND ORBITALS • There are regions with in the atom where we are likely to find electrons. You can think of these regions as clouds of electrons(aka electron density). • The denser the cloud… the more likely you will be to find an electron and vice versa.

21. The regions are described as orbitals. “An atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found.”Orbital ≠ orbit. The orbital is where the electron can be found 90% of the time.

22. There are 4 different kinds of orbital- all with different shapes • s- spherical in shape • p- Dumbbell shaped • d- An odd, complex shape • f- An odd, complex shape • ** The amount of energy an electron has determines which orbital it will occupy.**

23. ORBITALS AND ENERGY • Just as Bohr suggested the orbitals are assigned quantum numbers (n) and increase in energy level the farther away from the nucleus they get. However, each of the energy levels is subdivided in to our s, p, d and f orbitals. Not all energy levels have every type of orbital. Here is the trend: • Energy level n=1: 1s • E.L. n=2: 2s 2p • E.L. n=3: 3s 3p 3d • E.L. n=4: 4s 4p 4d 4f

24. ELECTRON SPIN • Electrons behave as though they spin on an axis and can either spin CLOCKWISE or COUNTERCLOCKWISE. The spinning of the electron creates a magnetic field. If it spins clockwise- the north end of the magnetic field points UP. If it spins counterclockwise- the magnetic north points DOWN.

25. The Pauli Exclusion Principle • In 1925, Austrian physicist Wolfgang Pauli used his exclusion principle to explain WHY spin is important to the arrangement of electrons in the atom. ***Each orbital can hold 2 electrons only and they must be of opposite spin.*** This is important because it tells us how many electrons are in each orbital and, also, each energy level.

26. Where do the Electrons GO? • Although all these orbitals are different, each can only hold a max of 2 electrons.

27. 5 Electronic Configurations • Electronic Configuration- the distribution of electrons within the orbitals of an atom. • “Electron configuration of atoms is determined by distributing the atom’s electrons among levels, sublevels, and orbitals based on a set of stated principles.”

28. DETERMINING ELECTRON CONFIGURATIONS • The Aufbau principle • Electrons are added ONE at a time. • Electrons always populate the lowest energy orbital available first- the fill in until all are accounted for. • The Pauli Exclusion Principle • An orbital can only hold a max of 2 electrons of opposite spin (a pair). It is possible to have only 1 electron (unpaired). • Hund’s Rule • Electrons fill equal-energy orbital with unpaired electrons(1 e- per orbital)first… then pair the electrons. EX. O • Sublevel notation (orbital notation) O 1s2 2s2 2p4 • The superscripts represent the # of electrons in the orbital. The #s in front of the orbital type tell you the energy level.

30. EXCEPTIONS TO THE AUFBAU PRINCIPLE • There are only 3 types of exceptions: • Cr [Ar]4s1 3d5 • Cu [Ar]4s1 3d10 • Nb [Kr]5s1 4d4 • Mo [Kr]5s1 4d5   • Ru [Kr]5s1 4d7 • Rh [Kr]5s1 4d8 • Pd [Kr]5s0 4d10 • Ag [Kr]5s1 4d10