1. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 1 Introductory Chemistry: A Foundation FIFTH EDITION by Steven S. Zumdahl University of Illinois

2. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 2 Modern Atomic Theory Chapter 10

3. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 3 Rutherford’s Atom Rutherford showed: –Atomic nucleus is composed of protons (positive) and neutrons (neutral). –The nucleus is very small compared to the size of the entire atom. Questions left unanswered: –How are electrons arranged and how do they move? The concept of a nuclear atom (charged electrons moving around the nucleus) resulted from Ernest Rutherford’s experiments.

4. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 4 The Rutherford Atom

5. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 5 Problems with Rutherford’s Nuclear Model of the Atom Electrons are moving charged particles Moving charged particles give off energy Therefore the atom should constantly be giving off energy And the electrons should crash into the nucleus and the atom collapse!!

6. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 6 Electromagnetic Radiation Classical physics says matter made up of particles, energy travels in waves Electromagnetic Radiation is radiant energy, both visible and invisible Electromagnetic radiation travels in waves All waves are characterized by: - velocity - wavelength, - frequency - amplitude

7. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 7 Electromagnetic Waves Velocity = c = speed of light –2.997925 x 10 8 m/s –all types of light energy travel at the same speed Amplitude = A = measure of the intensity of the wave, “brightness” Wavelength =  = distance between two consecutive peaks or troughs in a wave –generally measured in nanometers (1 nm = 10 -9 m) –same distance for troughs Frequency = = the number of waves that pass a point in space in one second –generally measured in Hertz (Hz), –1 Hz = 1 wave/sec = 1 sec -1 c = x

8. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 8 Types of Electromagnetic Radiation

9. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 9 Types of Electromagnetic Radiation Radiowaves = > 0.01 m, low frequency and energy Microwaves = 10 -4 m < < 10 -2 m Infrared (IR) –far = 10 -4 < < 10 -5 m –middle = 10 -5 < < 2 x 10 -6 m –near = 2 x 10 -6 < < 8 x 10 -7 m Visible = 8 x 10 -7 < < 4 x 10 -7 m –ROYGBIV Ultraviolet (UV) –near = 4 x 10 -7 < < 2 x 10 -7 m –far = 2 x 10 -7 < < 1 x 10 -8 m X-rays = 10 -8 < < 10 -10 m Gamma rays = < 10 -10

10. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 10 Planck’s Revelation Showed that light energy could be thought of as particles for certain applications Stated that light came in particles called photons Particles of light have fixed amounts of energy –Basis of quantum theory

11. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 11 Electromagnetic Radiation Pictured as Wave and Photons

12. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 12 The energy of the photon is directly proportional to the frequency of light Higher frequency = More energy in photons

13. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 13 Emission of Energy by Atoms Salts containing Li+, Cu2+, and Na+ dissolved in methyl alcohol set on fire: Li+ red Cu2+ green Na+ yellow

14. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 14 Emission of Energy by Atoms An excited lithium atom emitting a photon of red light to drop to a lower energy state.

15. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 15 Emission of Energy by Atoms/Atomic Spectra Atoms which have gained extra energy release that energy in the form of light The light atoms give off or gain is of very specific wavelengths called a line spectrum –light given off = emission spectrum –light energy gained = absorption spectrum –extends to all regions of the electromagnetic spectrum Each element has its own line spectrum which can be used to identify it

16. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 16 Emission Spectrum of Hydrogen

17. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 17 Atomic Spectra The line spectrum must be related to energy transitions in the atom. –Absorption = atom gaining energy –Emission = atom releasing energy Since all samples of an element give the exact same pattern of lines, every atom of that element must have only certain, identical energy states The atom is quantized –If the atom could have all possible energies, then the result would be a continuous spectrum instead of lines

18. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 18 The Energy Levels of Hydrogen

19. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 19 Hydrogen Atom (a) A sample of H atoms receives energy from an external source (b) The excited atoms (H) can release the excess energy by emitting photons

20. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 20 Quantized Nature of Energy The atom is quantized: 1.Only certain energy levels are allowed 2. If the atom could have all possible energies, then the result would be a continuous spectrum instead of lines

21. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 21 The Bohr Model of the Atom Explained spectra of hydrogen Energy of atom is related to the distance electron is from the nucleus Energy of the atom is quantized –atom can only have certain specific energy states called quantum levels or energy levels –when atom gains energy, electron “moves” to a higher quantum level –when atom loses energy, electron “moves” to a lower energy level –lines in spectrum correspond to the difference in energy between levels

22. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 22 Bohr’s Model Ground state- minimum energy levels –therefore atoms do not crash into the nucleus Ground state of hydrogen corresponds to having its one electron in an energy level that is closest to the nucleus Excited states- higher energy levels –the farther the energy level is from the nucleus, the higher its energy To put an electron in an excited state requires the addition of energy to the atom; bringing the electron back to the ground state releases energy in the form of light

23. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 23 Bohr’s Model

24. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 24 Problems with the Bohr Model Only explains hydrogen atom spectrum –and other 1 electron systems Neglects interactions between electrons Assumes circular or elliptical orbits for electrons - which is not true

25. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 25 Wave Mechanical Model of the Atom Experiments later showed that electrons could be treated as waves –just as light energy could be treated as particles –Louis de Broglie The quantum mechanical model treats electrons as waves and uses wave mathematics to calculate probability densities of finding the electron in a particular region in the atom –Schrödinger Wave Equation –can only be solved for simple systems, but approximated for others

26. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 26 Orbital-the probability map: describes the H electron in lowest energy state

27. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 27 Orbitals Solutions to the wave equation give regions in space of high probability for finding the electron - these are called orbitals –usually use 90% probability to set the limit –three-dimensional

28. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 28 Hydrogen’s 1s Orbital The electron can be found inside the sphere 90% of the time

29. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 29 Energy Levels: Principal energy Level Principal energy levels (n) - the discrete energy levels: Identify how much energy the electrons in the orbital have

30. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 30 Energy Levels: Sublevels Each principal energy level contains one or more sublevels –there are n sublevels in each principal energy level –each type of sublevel has a different shape and energy –s < p < d < f Each sublevel contains one or more orbitals –s = 1 orbital, p = 3, d = 5, f = 7

31. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 31 Principal levels divided into sublevels:

32. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 32 Orbital Labels: 1. The number tells the principal level 2. The letter tells the shape: s- sphere, p-two lobes

33. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 33 Principal Levels 1 and 2-shape of the orbitals that compose the sublevels:

34. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 34 3d orbitals - shapes and labels:

35. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 35 Wave Mechanical Model - Pauli Exclusion Principle No orbital may have more than 2 electrons Electrons in the same orbital must have opposite spins s sublevel holds 2 electrons p sublevel holds 6 electrons d sublevel holds 10 electrons f sublevel holds 14 electrons

36. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 36 for a many electron atom, build-up the energy levels, filling each orbital in succession by energy ground state 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p degenerate orbitals are orbitals with the same energy –each p sublevel has 3 degenerate p orbitals –each d sublevel has 5 degenerate d orbitals –each f sublevel has 7 degenerate f orbitals Orbitals, Sublevels & Electrons

37. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 37 Hund’s Rule for a set of degenerate orbitals, half fill each orbital first before pairing highest energy level called the valence shell –electrons in the valence shell called valence electrons –electrons not in the valence shell are called core electrons –often use symbol of previous noble gas to represent core electrons 1s 2 2s 2 2p 6 = [Ne]

38. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 38 Electron Configuration and the Periodic Table Elements in the same column on the Periodic Table have –Similar chemical and physical properties –Similar valence shell electron configurations Same numbers of valence electrons Same orbital types Different energy levels

39. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 39 s1s1 s2s2 d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 p 1 p 2 p 3 p 4 p 5 s2s2 p6p6 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 12345671234567

40. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 40 Writing Orbital Diagrams Magnesium (Mg): Z = 12 (12 electrons) Electron configuration: 1s 2 2s 2 2p 6 3s 2 Orbital Diagram: 1s 2s 2p 3s

41. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 41 Determining Electron Configuration Sulfur (S): Z = 16 (16 electrons) Electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 4 or [Ne] 3s 2 3p 4

42. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 42 The Modern Periodic Table Columns are called Groups or Families Rows are called Periods –Each period shows the pattern of properties repeated in the next period Main Groups = Representative Elements Transition Elements Bottom rows = Lanthanides and Actinides –really belong in Period 6 & 7

43. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 43 Atomic properties and the Periodic Table-Metallic Character Metals –malleable & ductile –shiny, lustrous –conduct heat and electricity –most oxides basic and ionic –form cations in solution –lose electrons in reactions - oxidized Nonmetals brittle in solid state dull electrical and thermal insulators most oxides are acidic and molecular form anions and polyatomic anions gain electrons in reactions - reduced Metalloids Also known as semi-metals Show some metal and some nonmetal properties

44. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 44 Metallic Character Metals are found on the left of the table, nonmetals on the right, and metalloids in between Most metallic element always to the left of the Period, least metallic to the right, and 1 or 2 metalloids are in the middle Most metallic element always at the bottom of a column, least metallic on the top, and 1 or 2 metalloids are in the middle of columns 4A, 5A, and 6A

45. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 45 Reactivity Reactivity of metals increases to the left on the Period and down in the column –follows ease of losing an electron Reactivity of nonmetals (excluding the noble gases) increases to the right on the Period and up in the column

46. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 46 Ionization Energy Minimum energy needed to remove a valence electron from an atom The lower the ionization energy, the easier it is to remove the electron Ionization Energy decreases down the group –valence electron farther from nucleus Ionization Energy increases across the period –left to right

47. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 47 Trends in Ionization Energy

48. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 48 Atomic Size Increases down column –valence shell farther from nucleus Decreases across period –left to right –adding electrons to same valence shell –valence shell held closer because more protons in nucleus

49. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 49 Trend in Atomic Size

50. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 50 Bohr’s Model Distances between energy levels decreases as the energy increases –light given off in a transition from the second energy level to the first has a higher energy than light given off in a transition from the third to the second, etc. –Electrons “orbit” the nucleus much like planets orbiting the sun 1st energy level can hold 2e -1, the 2nd 8e -1, the 3rd 18e -1, etc. –farther from nucleus = more space = less repulsion The highest energy occupied ground state orbit is called the valence shell

51. Copyright©2004 by Houghton Mifflin Company. All rights reserved. 51 Orbitals Solutions to the wave equation give regions in space of high probability for finding the electron - these are called orbitals –usually use 90% probability to set the limit –three-dimensional Orbitals are defined by three integer terms that are added to the wave equation to quantize it - these are called the quantum numbers Each electron also has a FIFTH quantum number to represent the direction of spin