Atomic Structure and Periodicity

1. Atomic Structure and Periodicity Chapter 7

2. Section 7.1 Electromagnetic Radiation • The electromagnetic spectrum organizes waves by wavelength, frequency, and energy l – wavelength- the length from a point on a wave to a corresponding point later in the wave. n – frequency- the number of times a full wave cycle passes by a reference point in one second.

3. The electromagnetic spectrum

4. Relationships between l, n, and E • Wavelength and frequency are inversely proportional (a long wave will take longer to pass by a reference point, thus making its frequency lower). • Energy is directly related to frequency. Higher frequency waves have more energy than lower frequency waves. • The speed of a wave is constant in a vacuum. (3.0 x 108 m/s = c, speed of light) l * n = c

5. Show Me Problem • A red light emits light of about 650 nm wavelength. What is the frequency of the red light? 650 nm  6.5 x 10-7m (6.5x10-7)(n) = (3.0 x 108) n = 4.61 x 1014 Hz l * n = c

6. Section 7.2 The Nature Of Matter • Max Plank • Quantizes energy • Matter can absorb energy, but only in whole number ratios of the term hn. Plank’s constant 6.626 x 10-34 J*s EQUATION: DE = (h*n)

7. Einstein’s Contribution • Albert Einstein • Quantizes Radiation • E = mc2 … Energy has mass and velocity. Electromagnetic Radiation must be made up of particles called photons. Duality of Light: Electromagnetic radiation has the capacity to behave both as a wave and a particle.

8. DeBroglie’s Contribution • Louis DeBroglie • Determines the Duality of Matter • If waves act as particles, do particles act as waves? • Set the Einstein and the Plank equations equal to each other. mc2 = E = hc/l m = h/cl l = h/mn Duality of Matter: Electrons have the capacity to behave both as a particle and a wave.

9. How can we be certain? • X Ray Crystallography Different color and shading patters appear as the electron “waves” cause diffraction: constructive and destructive interference with the x-rays that are exposed to the crystal. X Ray Diffraction Pattern of 2-terphenyl-4-yl-5-phenyl thiophene (PPPTP) X-Ray Diffraction Pattern of Beryl.

10. Atomic Spectrum of Hydrogen • Extensively studied by atomic theorists such as Bohr. • High energy sparks cause hydrogen gas molecules (H-H) to break apart suddenly, with some electrons in higher energy levels than would be expected normally. • As the electrons fall back to their ground states, energy is released. • Each color in the spectrum relates to an electron in a different energy level. • Plank’s equation can be used to determine the color of the light produced or the energy of the electron that is being observed.

11. More on the atomic spectrum of hydrogen

12. Section 7.4 The Bohr Model • Based upon the study of the Hydrogen Spectrum, Bohr designs paths for electrons to travel while orbiting the nucleus. [ORBITS] • Each orbit corresponded to a different energy level. energy of e- energy level nuclear charge (protons) Dr Quantum Video

13. Show Me • An electron in a hydrogen atom moves from energy level one to energy level 2. What is the change in energy the electron experiences? E1 = -2.178 x 10-18 (12/12)= -2.178 x 10-18 E2 = -2.178 x 10-18 (12/22)= -5.445 x 10-19 DE = E2-E1 = -5.445 x 10-19-(-2.178 x 10-18) DE = -1.633x10-18 Joules

14. Section 7.5 The Quantum Mechanical Model of the Atom • Determine that if the electron acts as a standing wave (a wave that is fixed in place), then there are only certain orientations for it to exist without causing destructive interference with itself. Heisenberg De Broglie Scrhödinger

15. Heisenberg Uncertainty Principle • There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle. • The more certain you are of the location of an electron, the less certain you can be of its momentum • The more certain you are of the momentum of an electron, the less certain you can be of its position. Dx * Dmn = h/4p

16. The Wave Equation • Schrodinger’s wave equation is used to define the location of electrons as waves. • A wave function is called an orbital. • ORBITALS = ORBITS • Wave functions are impossible to visualize. We picture the electron density map (aka: electron probability diagram) HY = EY

17. Section 7.6: Quantum Numbers Quantum numbers describe the properties of orbitals.

19. Common Orbital Shapes

20. Section 7.7 Orbital Shapes • Areas of high probability are separate by areas of low probability. (NODES) • Degenerate orbitals have different orientation or shape but the same ENERGY. • Lowest available energy level for an electron = ground state • Higher energy levels than expected = excited states

21. Section 7.8 Electron Spin and Pauli Principle • Electrons exhibit a fourth quantum number. • Electron spin quantum number (ms) • Values of ½ and -½. Indicates magnetic moment of electron. Electrons can only spin in one of 2 opposite directions. Pauli exclusion Principle: In a given atom, no two electrons Can have the same set of Quantum numbers.

22. Electron Configurations “Diagonal Rule” 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s

23. 3 Ways for Electron Configurations Electron Configuration Diagrams Nobel Gas Configurations Long-Hand Configurations

24. Electron Configuration Rule Summary • Electrons enter lowest energy orbitals first. • Only two electrons per degenerate orbital. • Electrons spread out among degenerate orbitals before pairing.

25. Exceptions to the Configuration Rules • A fully filled orbital is more stable than a partially filled orbital. • half-filled orbital is more stable than a more/less partially filled orbital. Cr Cu Eu Am Mo Ag

26. Cu expected configuration: 1s22s22p63s23p64s23d9 Cu actual configuration: 1s22s22p63s23p64s13d10 Cr expected configuration: 1s22s22p63s23p64s23d4 Cr actual configuration: 1s22s22p63s23p64s13d5 Copper and Chromium

27. Mo expected configuration: 1s22s22p63s23p64s23d104p65s24d4 Mo actual configuration: 1s22s22p63s23p64s23d104p65s14d5 Ag expected configuration: 1s22s22p63s23p64s23d104p65s24d9 Ag actual configuration: 1s22s22p63s23p64s23d104p65s14d10 Molybdenum and Silver

28. Eu expected configuration: 1s22s22p63s23p64s23d104p65s24d105p66s24f6 Eu actual configuration: 1s22s22p63s23p64s23d104p65s24d105p66s14f7 Am expected configuration: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d107s25f6 Am actual configuration: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d107s15f7 Europium and Americium

29. Electron Configuration Diagrams

30. Noble Gas Notation Expanded Titanium : 1s22s22p63s23p64s23d2 Noble Gas Titanium: [Ar] 4s23d2 Use the noble gas that comes before the element as a benchmark, then tack on the extra occupied orbitals.

31. Periodic Trends: Electron Configurations • Counting down tells what energy orbital. • Counting over tells how many electrons.

32. Periodic Trends Activity • http://academic.pgcc.edu/~ssinex/excelets/PT_interactive.xls Go to this excel sheet and click on the bottom tab labeled “atom properties”

33. Periodic Trends: Atomic Size Increasing Atomic Size

34. Periodic Trends: Ionization Energy Increasing 1st Ionization Energy

35. Periodic Trends: Electron Affinity Increasing electron affinity

36. Ion Size • Negative ions indicate a gain of electrons. • They are larger than the atom from whence they are formed. • Positive ions indicate a loss of electrons. • They are smaller than the atom from whence they are formed.