2. C HAPTER 5 A RRANGEMENT OF ELECTRONS IN ATOMS

3.  *Rutherford's model of the atom does not explain how the electrons fill the space **See Gold Foil experiment pg. 108  Evidence about the configuration of electrons in the orbitals (electron cloud) came from studying light:  Light (electromagnetic radiation) has a dual nature, meaning it behaves like a wave and a particle. 

4. W AVE DESCRIPTION OF LIGHT – o 1800's scientists believed that light was a beam of energy moving through space in the form of waves (like waves on a lake when a pebble is thrown in)

5. *A LL WAVES HAVE 4 CHARACTERISTICS : SEE PAGE amplitude - height of wave origin to crest wavelength (λ)- distance between crests Light is measured in nanometers (nm) frequency (v) how fast up and down (oscillations) units: waves/sec, Hertz (Hz), s -1 speed (c) - constant 2.998 x 10 8 m/s

6. C = SPEED OF LIGHT ( LATIN CELERATA ) Formula: c = λv* *wavelength and frequency are inversely proportional ** meaning that if wavelength decreases then frequency increases & vice versa.

7. #1 E XAMPLE PROBLEM : What is the frequency of light that has a wavelength of 450 nm? hint: convert nm to m (1m = 1 x 10 9 nm)

8. #2 E XAMPLE PROBLEM : What is the wavelength of electromagnetic radiation if its frequency is 4.5 x 10 -3 Hz?

9. E XIT Q UESTION : 5 POINTS Write down 3 things that you learned today. Write down one thing you don’t understand.

10. P ARTICLE DESCRIPTION OF LIGHT 1900's experiments showed that light behaved like a stream of extremely tiny, fast moving particles.

11. E VIDENCE THAT SUPPORTS PARTICLE BEHAVIOR 1) photoelectric effect - refers to the emission of electrons from a metal when light shines on the metal (but only if the frequency was at a certain minimum) ex/ solar powered items work if you have enough light

12. M ORE EVIDENCE FOR PARTICLE BEHAVIOR 2) Max Planck - studied light emitted from hot metal objects (like a hot horseshoe glows). He suggests that objects emit energy in small specific amounts called quanta.

13. Quantum - minimum quantity of energy that can be lost or gained by an atom To calculate the energy of a quantum of light use formula: E = hv Where: E = energy (in Joules units) h = 6.626 x 10 -34 Js (Joule seconds) Planck's constant v = frequency

14. A LBERT E INSTEIN (1905) Introduces the wave-particle dual nature of light. wave & particle behavior each particle carries a quantum of energy. EM radiation is absorbed by matter in whole numbers of photons.

15. photon - particle of light (EM radiation) having zero mass and carrying a quantum of energy. E photon = hv

16. E XAMPLE PROBLEM : Using: E photon = hv Calculate the frequency for a photon of light that has an energy 3.2 x 10 -19 J.

17. H YDROGEN ’ S LINE EMISSION SPECTRUM Niels Bohr passed electric current through hydrogen gas PINK colored light emitted When energy is added to an atom, electrons become excited & move to higher energy level.

18. A photon is emitted when the electrons move back to a more stable, GROUND state. Ground state – lowest energy state of an atom Excited state – state in which an atom has a higher potential energy than its ground state. E photon = E 2 – E 1 = hv

19. The energy of this photon is equal to the difference in energy between the atom’s initial state and its final state.

20. B OHR M ODEL OF THE H YDROGEN ATOM 1913 Bohr links the photon emission of hydrogen to a model of the atom’s electron. See p. 129 Electron circles in orbits (defined paths) Electron has a fixed energy Each concentric circle orbit had an empty space in between where the electron could not exist (ladder analogy p. 129)

21. Explanation of the spectral lines produced by hydrogen: An electron cannot gain or lose energy. It can move to a higher energy orbit by gaining an amount of energy equal to the difference in final and initial states.

22. L OUIS D E B ROGLIE (“ DE BROYLEE ”) 1924 He proposed an equation that suggested that any matter with mass and velocity has a corresponding wavelength.

23. Setting both energy equations equal to each other: E = mc 2 E=h v mc 2 = h v (substitute v with wavelength from c = λ v) Wavelength(λ) = h/mc

24. W ERNER H EISENBERG 1927 e - s are detected by their interaction with photons. This interaction will change both the direction and position of the e -. Heisenberg uncertainty principle States: It is impossible to determine simultaneously both position and velocity of an e -

25. Heisenberg uncertainty principle States: It is impossible to determine simultaneously both position and velocity of an e -

26. Therefore, e - s are located in orbitals or 3-D clouds of probable location (not neat orbits like Bohr’s model nor Rutherford’s planetary model)

27. Erwin Schrodinger came up with an equation that treated electrons in atoms as waves. Quantization of electron energies was an outcome of his equation (vs. Bohr’s theory that assumed quantization as a fact)

28. S EC 1 1. For electromagnetic radiation, c (speed of light) equals _________________________. 2. A quantum of electromagnetic energy is called _______________. 3. The energy of a photon is related to its _____________. 4. If electrons in an atom have the lowest possible energies, the atom is in the ________________. 5. Bohr’s theory helped explain why excited hydrogen gas gives off certain ___________ of light. 6. According to Bohr’s theory, an excited atom would _______________ energy.

29. S ECTION 2 R EVIEW Q’ S 1. A three-dimensional region around a nucleus where an electron may be found is called a(n) ____________. 2. Unlike in an orbit, in an orbital an electron’s position cannot be known _______________. 3. What are the 4 quantum numbers and what do they represent? 4. What are the shapes of the orbitals? 5. How many electrons fit in each orbital? 6. What is the difference between a 2s orbital and a 4s orbital?

30. S EC 2 1. How many orbital shapes are possible at the 2 nd energy level? 3 rd energy level? 2. An electron for which n= 5 has more _____ than an electron for which n=3. 3. If 8 electrons completely fill a main energy level, what is n?

31. S ECTION 3 R EVIEW Q’ S 1. Draw the diagonal rule. What does this rule show? 2. Know the 3 rules for writing electron configurtions. 3. Write the electron configuration for Si. 4. Draw the orbital diagram for Mg. 5. What element has the following configuration: 1s 2 2s 2 2p 6 3s 1 ? 6. How many electrons in the highest energy level of a bromine atom? 7. Which element has the electron configuration of [Ar]4s 2 3d 10 4p 5