Chapter 7: Periodicity and Atomic Structure

1. Chapter 7: Periodicity and Atomic Structure Renee Y. Becker Valencia Community College CHM 1045

2. Light and Electromagnetic Spectrum • Several types of electromagnetic radiation make up the electromagnetic spectrum

3. Light and Electromagnetic Spectrum Frequency, : The number of wave peaks that pass a given point per unit time (1/s) Wavelength, : The distance from one wave peak to the next (nm or m) Amplitude: Height of wave Wavelength x Frequency = Speed (m) x (s-1) = c (m/s)

4. Light and Electromagnetic Spectrum

5. The Planck Equation E = h  E = hc /  h = Planck’s constant, 6.626 x 10-34 J s 1 J = 1 kg m2/s2

6. Example1: Light and Electromagnetic Spectrum • The red light in a laser pointer comes from a diode laser that has a wavelength of about 630 nm. What is the frequency of the light? c = 3 x 108 m/s

7. Atomic Spectra • Atomic spectra: Result from excited atoms emitting light. • Line spectra: Result from electron transitions between specific energy levels. • Blackbody radiation is the visible glow that solid objects emit when heated.

8. Atomic Spectra • Max Planck (1858–1947): proposed the energy is only emitted in discrete packets called quanta. The amount of energy depends on the frequency: E = energy = frequency  = wavelength c = speed of light h = planck’s constant hc 34 - E h h 6 . 626 10 J s = n = = ´ × l

9. Atomic Spectra Albert Einstein (1879–1955): Used the idea of quanta to explain the photoelectric effect. • He proposed that light behaves as a stream of particles called photons • A photon’s energy must exceed a minimum threshold for electrons to be ejected. • Energy of a photon depends only on the frequency. E = h 

10. Atomic Spectra

11. Example 2: Atomic Spectra • For red light with a wavelength of about 630 nm, what is the energy of a single photon and one mole of photons?

12. h h : For Light l = = mc p h h For a Particle : l = = mv p Wave–Particle Duality • Louis de Broglie (1892–1987): Suggested waves can behave as particles and particles can behave as waves. This is called wave–particle duality. m = mass in kg p = momentum (mc) or (mv)

13. Example 3: Wave–Particle Duality • How fast must an electron be moving if it has a de Broglie wavelength of 550 nm? me = 9.109 x 10–31 kg

14. Quantum Mechanics • Niels Bohr (1885–1962): Described atom as electrons circling around a nucleus and concluded that electrons have specific energy levels. • Erwin Schrödinger (1887–1961): Proposed quantum mechanical model of atom, which focuses on wavelike properties of electrons.

15. Quantum Mechanics • Werner Heisenberg (1901–1976): Showed that it is impossible to know (or measure) precisely both the position and velocity (or the momentum) at the same time. • The simple act of “seeing” an electron would change its energy and therefore its position.

16. Quantum Mechanics • Erwin Schrödinger (1887–1961): Developed a compromise which calculates both the energy of an electron and the probability of finding an electron at any point in the molecule. • This is accomplished by solving the Schrödinger equation, resulting in the wave function

17. Quantum Numbers • Wave functions describe the behavior of electrons. • Each wave function contains four variables called quantum numbers: • Principal Quantum Number (n) • Angular-Momentum Quantum Number (l) • Magnetic Quantum Number (ml) • Spin Quantum Number (ms)

18. Quantum Numbers • Principal Quantum Number (n): Defines the size and energy level of the orbital. n = 1, 2, 3,  • As n increases, the electrons get farther from the nucleus. • As n increases, the electrons’ energy increases. • Each value of n is generally called a shell.

19. Quantum Numbers • Angular-Momentum Quantum Number (l): Defines the three-dimensional shape of the orbital. • For an orbital of principal quantum number n, the value of l can have an integer value from 0 to n – 1. • This gives the subshell notation:l = 0 = s orbital l = 3 = f orbital l = 1 = p orbital l = 4 = g orbital l = 2 = d orbital

20. Quantum Numbers • Magnetic Quantum Number (ml): Defines the spatial orientation of the orbital. • For orbital of angular-momentum quantum number, l, the value of ml has integer values from –l to +l. • This gives a spatial orientation of:l = 0 giving ml = 0 l = 1 giving ml = –1, 0, +1l = 2 giving ml = –2, –1, 0, 1, 2, and so on…...

21. Quantum Numbers • Magnetic Quantum Number (ml): –l to +l S orbital 0 P orbital -1 0 1 D orbital -2 -1 0 1 2 F orbital -3 -2 -1 0 1 2 3

22. Quantum Numbers • Spin Quantum Number: ms • The Pauli Exclusion Principle states that no two electrons can have the same four quantum numbers.

23. Quantum Numbers

24. Example 4: Quantum Numbers • Why can’t an electron have the following quantum numbers? (a) n = 2, l = 2, ml = 1 (b) n = 3, l = 0, ml = 3 (c) n = 5, l = –2, ml= 1

25. Example 5: Quantum Numbers • Give orbital notations for electrons with the following quantum numbers: • n = 2, l = 1 (b) n = 4, l = 3 (c) n = 3, l = 2

26. Electron Radial Distribution • s Orbital Shapes: Holds 2 electrons

27. Electron Radial Distribution • p Orbital Shapes: Holds 6 electrons, degenerate

28. Electron Radial Distribution • d and f Orbital Shapes: d holds 10 electrons and f holds 14 electrons, degenerate

29. Effective Nuclear Charge • Electron shielding leads to energy differences among orbitals within a shell. • Net nuclear charge felt by an electron is called the effective nuclear charge (Zeff). • Zeff is lower than actual nuclear charge. • Zeff increases toward nucleus ns > np > nd > nf

30. Effective Nuclear Charge