Electrons in Atoms and the Periodic Table Chapter 9

1. Electrons in Atomsand the Periodic TableChapter 9 Tro, 2nd ed.

2. A Brief History of Atomic Theory Greeks were the first to suggest that matter was made up of small particles called atoms Early chemists performed experiments to find these particles Their experiments led to Dalton’s Atomic Theory Dalton’s Atomic Theory was revised further to the Thompson and Rutherford “nuclear” models of the atom Further work has led to the modern nuclear atomic theory, which is still being revised

3. A Brief History of Atomic Theory ~1900 Max Planck: energy emitted in bursts called quanta ~1905 Einstein used quanta to explain properties of light Called quantum of light a photon ~1911 Rutherford proposed “nuclear atom” Nucleus is a tiny center in the atom: Contains most of the mass Contains all of the positive charge (Thought electrons in orbit like mini-solar system - not!)

4. WHAT YOU NEED TO KNOW ABOUT ENERGY AND LIGHT (& e-s & photons) Light is part of the electromagnetic spectrum of radiation Continuous spectrumis likea “rainbow” Discrete spectrum is lines of specific colors of light Each line of color has its own energy associated with it Light in a vacuum moves at constant speed c = 2.998 X 108 m/sec (memorize!)

5. WHAT YOU NEED TO KNOW ABOUT ENERGY AND LIGHT (& e-s & photons) Electromagnetic Radiation - Examples light from the sun x-rays microwaves radio waves television waves radiant heat All show wavelike behavior. Each travels at the same speed in a vacuum: 2.998 x 108 m/s

6. X-rays are part of the electromagnetic spectrum visible light is part of the electromagnetic spectrum Infrared light is part of the electromagnetic spectrum The Electromagnetic Spectrum Know visible light as ROYGBIV.

7. Characteristics of a Wave Light has the properties of a wave. Wavelength, l (measured from trough to trough) Wavelength, l (measured from peak to peak)

8. Characteristics of a Wave Frequency, n, is the number of wavelengths that pass a particular point per second.

9. Characteristics of a Wave Speed is how fast a wave moves through space.

10. WHAT YOU NEED TO KNOW ABOUT ENERGY AND LIGHT (& e-s & photons) Each specific “color” of visible light has a specific wavelength called land a specific frequency called  and both relate to speed of light c = l Wavelength, l, is length unit: meters/wavelength Frequency, n, is number of wavelengths/second Speed of light: c = meters * wavelengths = m/s wavelength second Energy (Joules) of one photon of light of specific wavelength or frequency: E = hn = Joules/photon h = Planck’s constant (6.626 x 10-34 J.s)

11. The Bohr Atom At high temperatures or voltages, elements in the gaseous state emit light of different colors. When the light is passed through a prism or diffraction grating a line spectrum results. Niels Bohr, a Danish physicist, in 1912-1913 carried out researchon the hydrogen atom.

12. Each element has its own unique set of spectral emission lines that distinguish it from other elements. These colored lines indicate that light is being emitted only at certain wavelengths. Line spectrum of hydrogen. Each line corresponds to the wavelength of the energy emitted when the electron of a hydrogen atom, which has absorbed energy falls back to a lower principal energy level.

13. Tro: Figure 9.8: each element produces its own unique emission spectrum. Spectra

14. Energy is absorbed by e-, then emitted as a photon of light. Tro: Figure 9.11

15. The BohrAtom E1 E2 E3 An electron can have one of several possible energies depending on its orbit.

16. An electron has a discrete energy when it occupies an orbit. The Bohr Atom Electrons revolve around the nucleus in orbits that are located at fixed distances from the nucleus.

17. Different lines of the hydrogen spectrum correspond to different electron energy level shifts. The color of the light emitted corresponds to one of the lines of the hydrogen spectrum. The Bohr Atom When an electron falls from a higher energy level to a lower energy level a quantum of energy in the form of light is emitted by the atom.

18. The BohrAtom Light is not emitted continuously, but is emitted in discrete packets called photons

19. CALCULATIONS FOR l, n AND E OF A PHOTON EMITTED BY AN ELECTRON When e- in H atom moves from energy level 1 up to energy level 4 and then drops back down to energy level 2, we see a photon of light emitted that has a wavelength, l, of 4.86 x 10-7 m. Calculate the frequency of the light and the energy of the photon emitted. c=ln Rearrange to get frequency, n n=c/l = (2.998x108 m/s)/(4.86x10-7 m) = 6.17 x 1014 /s E=hn = 6.626x10-34 J.s * 6.17 x 1014 /s = 4.09 x 10-19 J/photon of light

20. MODERN ATOMIC THEORY Bohr’s calculations succeeded very well for the hydrogen atom. Bohr’s methods did not succeed for heavier atoms. So…more theoretical work on atomic structure was needed.

21. MODERN ATOMIC THEORY Thompson had shown that light, which is photons of energy, had the properties of matter as well. In 1924, Louis De Broglie suggested that all matter must also have wave properties. De Broglie showed that the wavelength of ordinary sized objects, such as a baseball, are too small to be observed. For objects the size of an electron the wavelength can be detected.

22. MODERN ATOMIC THEORY In 1926, Schröedinger created a mathematical model that showed electrons as waves. Schröedinger’s work led to a new branch of physics called wave or quantum mechanics. Using Schröedinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined. The actual location of an electron within an atom cannot be determined (Heisenberg Uncertainty Principle).

23. MODERN ATOMIC THEORY Based on wave mechanics it is clear that electrons are not revolving around the nucleus in orbits. Instead of being located in orbits, the electrons are located in orbitals. An orbital is a region around the nucleus where there is a high probability of finding an electron.

24. MODERN ATOMIC THEORY According to Bohr the energies of electrons in an atom are quantized. The wave-mechanical model of the atom also predicts discrete principal energylevels within the atom.

25. Energy Levelsof Electrons Principal energy levels, n = 1,2,3...(also called principal quantum number) Each energy level contain(s) sublevel(s) called s, p, d, and f (the angular momentum quantum number, l) Within sublevels are orbitals (designated by orientation in 3-D space, called the magnetic quantum number, ml) Each orbital can hold 2 e-s max (each electron is assigned a spin quantum number, ms)

26. The first four principal energy levels of the hydrogen atom As n increases, the energy of the electron increases Each level is assigned a principal quantum number, n Energy levels in atoms

27. Energy levels in atoms The number of sublevels equals the assigned energy level. For n=1, there is one sublevel, s. For n=2, there are two sublevels, s & p, etc. The sublevels have the quantum designation, l. The maximum value is n -1. Each principal energy level is subdivided into sublevels.

28. Within sublevels the electrons are found in orbitals. An s orbital is round and soft, like a nerf ball. The shape represents the highest probability where the electron might be found.

29. An atomic orbital can hold a maximum of two electrons. An electron can spin in one of two possible directions represented by ↑ or ↓. The two electrons that occupy an atomic orbital must have opposite spins. This is known as the Pauli Exclusion Principal.

30. A p sublevel is made up of three p-type orbitals. Each p orbital has two lobes and can hold a maximum of two electrons. Since there are three orbitals, a p sublevel can hold a maximum of 6 electrons.

31. The three p orbitals all center at the atom’s nucleus… pz py px …and occupy one of the three axes of 3-D space.

32. A d sublevel is made up of five orbitals. The five d orbitals lie in different planes and point in different directions. Each d orbital can hold a maximum of two electrons. A d sublevel can hold a maximum of 10 electrons.

33. Tro Figure 9.19: The number of subshells within a shell is equal to the value of n, the principal quantum number. Shells & Subshells Etc, etc, etc: What subshells exist for n = 5?

34. Energy Levelsof Electrons Pauli Exclusion Principle: Each orbital can hold a max of 2 e-s, so possibilities are 0, 1 or 2 e-s. s has only 1 orbital  2 e-s MAX p has 3 orbitals  6 e-s MAX d has 5 orbitals  10 e-s MAX f has 7 orbitals  14 e-s MAX (g has ? Orbitals  ?? e-s MAX)

35. Energy Levelsof Electrons Max is 2 e-s because of a property of e-s called spin. (Pauli Exclusion Principle) Each e- is spinning on its axis like the earth. Any spinning charge creates a magnetic field, with N and S poles. The direction of spin determines which is North. The e-s will pair up in an orbital so their N poles are opposite each other.

36. Atomic Structure of the First 18 Elements: Use these guidelines The ground state of the electron is the lowest energy orbital it can occupy. Higher energy orbitals are excited states. The distribution of electrons into the various energy shells and subshells in an atom in its ground state is called its electron configuration. Each energy shell and subshell has a maximum number of electrons it can hold: s = 2, p = 6, d = 10, f = 14 Place electrons in the energy shells and subshells in order of energy, from low energy up: the Aufbau Principal.

37. Atomic Structure of the First 18 Elements: Use these guidelines No more than two electrons can occupy one orbital Electrons occupy the lowest energy orbitals available, the ground state. They enter a higher energy orbital only after the lower orbitals are filled. (Aufbau again.) For the atoms beyond hydrogen, orbital energies vary as “s<p<d<f” for a given value of n. Each orbital in a sublevel is occupied by a single electron before a second electron enters. For example, all three p orbitals must contain one electron before a second electron enters a p orbital. (Hund’s Rule)

38. 6d 7s 5f 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s Energy After 3p is filled, the next lowest energy is 4s, not 3d. See all the overlap beyond 3p.

39. Order of Subshell Fillingin Ground State Electron Configurations 1s 2s2p 3s3p3d 4s4p4d4f 5s5p5d5f 6s6p6d 7s Start by drawing a diagram putting each energy shell on a row and listing the subshells, (s, p, d, f), for that shell in order of energy, (left-to-right) Next, draw arrows through the diagonals, looping back to the next diagonal each time

40. Arrangement of electrons within their respective sublevels. 2p6 Writing Electron Configurations Number of electrons in sublevel orbitals Type of orbital Principal energy level

41. Order of Subshell Fillingin Ground State Electron Configurations Try using the drawing with the arrows to do aluminum, phosphorus and chlorine. Al 1s22s22p63s23p1 P 1s22s22p63s23p3 Cl 1s22s22p63s23p5

42. Orbital Box Notation In the following diagrams boxes represent orbitals. Electrons are indicated by arrows: ↑ or ↓ for spin. Electron configurations can be spdf or orbital box. See both as follows.

43. He Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins. Filling the 1s sublevel H ↑ 1s1 Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. ↑ ↓ 1s2

44. ↑ ↓ ↑ ↓ Be 1s 2s The 2s orbital fills upon the addition of beryllium’s third and fourth electrons. Filling the 2s sublevel ↑ 1s22s1 Li 1s 2s The 1s orbital is filled. Lithium’s third electron will enter the 2s orbital. ↑ ↓ 1s22s2

45. ↑ ↓ ↑ ↓ C 1s 2s 2p The second p electron of carbon enters a different p orbital than the first p electron so as to give carbon the lowest possible energy. ↑ ↓ ↑ ↓ N 1s 2s 2p The third p electron of nitrogen enters a different p orbital than its first two p electrons to give nitrogen the lowest possible energy. ↑ ↓ ↑ ↑ ↓ B 1s22s22p1 1s 2s 2p Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. ↑ ↑ 1s22s22p2 ↑ ↑ ↑ 1s22s22p3

46. ↑ ↓ ↑ ↓ O 1s 2s 2p There are four electrons in the 2p sublevel of oxygen. One 2p orbitals is has a second electron, with a spin opposite to the electron already in the orbital. ↑ ↓ ↑ ↓ F 1s 2s 2p There are five electrons in the 2p sublevel of fluorine. Two of the 2p orbitals are now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. ↑ ↑ 1s22s22p4 ↑ ↓ ↑ ↓ ↑ ↓ ↑ 1s22s22p5

47. ↑ ↓ ↑ ↑ ↓ ↑ ↓ ↑ ↓ ↓ Ne 1s22s22p6 1s 2s 2p There are 6 electrons in the 2p sublevel of neon, which fills the sublevel.

48. ↑ ↓ ↑ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ Mg 3s 1s 2s 2p The 3s orbital fills upon the addition of magnesium’s twelfth electron. ↑ ↓ ↑ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ Na 1s22s22p63s1 3s 1s 2s 2p The 2s and 2p sublevels are filled. The next electron enters the 3s sublevel of sodium. ↓ 1s22s22p63s2

50. After element 18, the 4s sublevel is filled before the 3d is filled.