I. Waves & Particles (p. 134-145)

1. I. Waves & Particles(p. 134-145) Ch. 5- Electrons in Atoms Unit 7 Targets: The Electronic Structure of Atoms(Chap 5) • I CAN Utilize appropriate scientific vocabulary to explain scientific concepts. • I CAN Perform calculations involving the energy, wavelength and frequency of electromagnetic waves. • I CAN Perform calculations to determine the de Broglie wavelength of any object. • I CAN Compare and contrast Bohr’s solar system model with Schrodinger’s wave mechanical model. • I CAN Predict the movement of electrons from the ground state to an excited state, back to the ground state. • I CAN Recognize elements based on their emission spectra. (Flame lab). • I CAN Generate electron configurations for elements: full electron configuration, noble gas configuration, orbital diagram. • I CAN Differentiate relationships between orbitals, sublevels and energy levels. • I CAN Correlate the relationship between electrons in a single orbital or single sublevel.

2. A. Waves • Electromagnetic radiation- a form of energy that exhibits wavelike behavior. • Wavelength() - length of one complete wave • Frequency () - # of waves that pass a point during a certain time period • SI UNIT for frequency is hertz (Hz) = 1 wave / s • Amplitude (A) - distance from the origin to the trough or crest

3.  crest A A origin trough  A. Waves greater amplitude (intensity) greater frequency (color)

4. Wave Nature of Light (& particles) • To understand the electronic structure of atoms we must understand light and how it is emitted or absorbed by substances. • We will examine visible light a type of Electromagnetic Radiation (EM) which carries (radiant) energy through space (speed of light) and exhibits wavelike behavior. • Also need to think of light as particle, to help understand how EM radiation and atoms interact

5. B. EM Spectrum HIGH ENERGY LOW ENERGY

6. Characteristics of EM radiation • Move through a vacuum at the ‘speed of light’ 3.00 x 108 m/s • Behaves like waves that move through water, which are the result of a transfer of energy to the water (from a stone), expressed as up and down movement of water • Both electric and magnetic properties

7. R O Y G. B I V red orange yellow green blue indigo violet B. EM Spectrum HIGH ENERGY LOW ENERGY

8. Wave Speed = (distance between wave peaks) x (frequency) = (wavelength) x (frequency) EM radiation moves through a vacuum at the “speedof light” 3.00 x 108 m/s also calledc. • A lower energy wave (infrared and red) has a longer wavelength() and lower frequency(f)  • A higherenergy wave (blue - violet) has a shorter wavelength() and higher frequency(f).

9. B. EM Spectrum c =  c: speed of light (3.00  108 m/s) : wavelength (m, nm, etc.) : frequency (Hz) Frequency & wavelength are inversely proportional

10. WORK:  = c   = 3.00  108 m/s 4.34  10-7 m B. EM Spectrum GIVEN:  = ?  = 434 nm = 4.34  10-7 m c = 3.00  108 m/s = 6.91  1014 Hz EX: Find the frequency of a photon with a wavelength of 434 nm.

11. C. Quantum Theory • Planck (1900) • Observed - emission of light from hot objects • Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta) • Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom.

12. C. Quantum Theory • Planck proposed that the energy, E, of a single quantum energy packet equals a constant (h) times its frequency • The energy of a photon is proportional to its frequency. E = h E: energy (J, joules) h: Planck’s constant (6.6262  10-34 J·s) : frequency (Hz)

13. C. Quantum Theory GIVEN: E = ?  = 4.57  1014 Hz h =6.6262  10-34 J·s WORK: E = h E = (6.6262  10-34 J·s) (4.57  1014 Hz) E = 3.03  10-19 J EX: Find the energy of a red photon with a frequency of 4.57  1014 Hz.

14. C. Quantum Theory • Planck (1900) • Observed - emission of light from hot objects • Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta) • Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom.

15. Classical Theory Quantum Theory C. Quantum Theory vs. Planck (1900)

16. Quantum Theory • Energy is always emitted or absorbed in whole number multiples of hv, such as hv, 2 hv, 3 hv, 4hv, …. The allowed energies are quantized, that is their values are restricted to certain quantities.   • The notion of quantized rather than continuous energies is strange. Consider a ramp and a staircase, on a ramp you can vary the length your steps and energy used on the walk up. When walking up steps you must exert exactly the specific amount of energy needed to reach the next step. Your steps on steps are quantized, you cannot step between them.

17. C. Quantum Theory • Einstein (1905) • Observed - photoelectric effect

18. C. Quantum Theory • Einstein (1905) • Concluded - light has properties of both waves and particles (photons) “wave-particle duality” • Photon - particle of light that carries a quantum of energy • Used planck’s quantum theory to deduced that: Ephoton= hv

19. II. Bohr Model of the Atom

20. A. Line-Emission Spectrum • DEFINITION: A set of frequencies of EM waves emitted by atoms of the element. EX. neon light absorb electrical energy, e- get excited, become somewhat unstable and release energy in the form of light, a prism separates the light into an atomic emission spectrum. excited state PHOTON OUT ENERGY IN ground state

21. B. Bohr Model • e- exist only in orbits with specific amounts of energy called energy levels • Therefore… • e- can only gain or lose certain amounts of energy • only certain photons are produced

22. B. Bohr Model • Energy of photon depends on the difference in energy levels • Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 6 5 4 3 2 1

23. C. Other Elements Helium • Bohr’s calculations only worked for hydrogen!  • Each element has a unique bright-line emission spectrum. • “Atomic Fingerprint”

24. C. Other Elements • Examples: • Iron • Now, we can calculate for all elements and their electrons – next section 

25. III. Quantum Model of the Atom

26. VISIBLE LIGHT ELECTRONS A. Electrons as Waves EVIDENCE: DIFFRACTION PATTERNS • Louis de Broglie (1924) • Applied wave-particle theory to e- • e- exhibit wave properties

27. B. Quantum Mechanics • Heisenberg Uncertainty Principle • Impossible to know both the velocity and position of an electron at the same time

28. B. Quantum Mechanics • SchrödingerWave Equation (1926) • finite # of solutions  quantized energy levels • defines probability of finding an e- Wavelength = plank’s constant/mass X velocity

29. Radial Distribution Curve Orbital B. Quantum Mechanics • Orbital (“electron cloud”) • Region in space where there is 90% probability of finding an e-

30. UPPER LEVEL C. Quantum Numbers • Four Quantum Numbers: • Specify the “address” of each electron in an atom

31. C. Quantum Numbers 1. Principal Quantum Number ( n ) • Main energy level occupied the e- • Size of the orbital • n2 = # of orbitals in the energy level

32. s p d f C. Quantum Numbers 2. Angular Momentum Quantum # ( l ) • Energy sublevel • Shape of the orbital

33. C. Quantum Numbers • n = # of sublevels per level • n2 = # of orbitals per level • Sublevel sets: 1 s, 3 p, 5 d, 7 f

34. C. Quantum Numbers 3. Magnetic Quantum Number ( ml) • Orientation of orbital around the nucleus • Specifies the exact orbitalwithin each sublevel

35. C. Quantum Numbers px py pz

36. 2s 2px 2py 2pz C. Quantum Numbers Orbitals combine to form a spherical shape.

37. C. Quantum Numbers 4. Spin Quantum Number ( ms) • Electron spin  +½ or -½ • An orbital can hold 2 electrons that spin in opposite directions.

38. C. Quantum Numbers 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin # energy level sublevel (s,p,d,f) orientation electron • Pauli Exclusion Principle • No two electrons in an atom can have the same 4 quantum numbers. • Each e- has a unique “address”:

39. Feeling overwhelmed? Read Section 4-2!