Arrangement of Electrons in Atoms

1. Arrangement of Electrons in Atoms 4.1 The Development of a New Atomic Model 4.2 The Quantum Model of the Atom 4.3 Electronic Configurations

2. 4.1 The Development of a New Atomic Model • Properties of Light • The Wave Description of Light • Light is a form of electromagnetic radiation. Its energy is transmitted in the form of waves.

3. 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.

4. The electromagnetic spectrum spans a wide range of wavelengths, from radio to X-rays. • Wavelength (λ) – the distance between corresponding points on adjacent waves

5. Frequency (v)- the number of waves the pass a given point in a specific time, usually one second. Expressed in waves/second or hertz. Named after Heinrich Hertz. Frequency and Wavelength: c = λv c- speed of light

6. “Classic” Model? • 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 element give off different colors when burned?

7. 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.

8. E=hv • E= amount of energy, • v= the frequency • 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.

9. 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. Quanta were 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.

10. How much energy does each photon carry?? Planck’s equation tells us. Ephoton=hv • 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

11. 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!!!!

12. 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. • Continuous Spectrum: • Emission of a continuous range of frequencies of electromagnetic radiation

13. 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:

14. *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).

15. The H-Atom Line-Emission Spectrum • Ground state- The lowest energy state of an atom. • Excited state- higher potential energy than ground state.

16. *The spectral lines 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.

17. 4.2 The Quantum Model of the Atom • If light is a wave-icle… can electrons be wave-icles too? • The Frenchmen, Louis de Broglie first suggested and proved it possible mathematically in 1924. And just like Planck it took American ingenuity to prove him correct.

18. 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!

20. The Heisenberg 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 same 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.

21. The Schrödinger Wave Equation • 1926, Austrian physicist Erwin Schrödinger developed an equation were e- are treated like waves. • With Heisenberg, The Quantum Theory was developed: • Mathematical description of the wave properties of electrons & other very small particles.

22. **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. • Orbital- a 3-D region around the nucleus that indicates the probable location of a electron

23. Atomic Orbitals &Quantum Numbers • Quantum numbers- specify the properties of atomic orbitals & the properties of electrons in orbitals. • Principal Quantum Number- n, indicates the main energy level occupied by electron.

24. Angular Momentum Quantum Number- laka azimuthal quantum number gives the orbital angular momentum. In chemistry, this quantum number is very important, since it specifies the shape of an atomic orbital and strongly influences chemical bonds and bond angles. • l=0 is called a s orbital, l=1, a p orbital, l=2, a d orbital and l=3, a f orbital.

25. Magnetic quantum number- m, indicates the orientation of an orbital around the nucleus.

27. Spin quantum number- has only 2 possible values (-1/2 , +1/2) which indicate the 2 fundamental states of an electron in an orbital. • A single orbital can hold up to 2 electrons which must be of opposite spin.

28. In 1925, Austrian physicist Wolfgang Pauli 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

29. Although all these orbitals are different, each can only hold a max of 2 electrons.

30. 4.3 Electron 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.”

31. Rules Governing 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.

32. EXCEPTIONS TO THE AUFBAU PRINCIPLE • There are only 3 exceptions: • Cr [Ar]4s1 3d5 =All s and d subshells are half full • Cu [Ar]4s1 3d10 =Prefers a filled d subshell, leaving s with 1 • Nb [Kr]5s1 4d4 <=5s and 4d energy levels are close • Mo [Kr]5s1 4d5   similar to Cr above • Tc [Kr]5s2 4d5   (not special, but think of Hund's rule) • Ru [Kr]5s1 4d7 <= Only 1 5s electron • Rh [Kr]5s1 4d8 <= in both • Pd [Kr]5s0 4d10 <= Note filled 4d and empty 5s • Ag [Kr]5s1 4d10 <= partial filled 5s, but filled d