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Modern Atomic Theory Chapter 10

Modern Atomic Theory Chapter 10

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Modern Atomic Theory Chapter 10

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  1. Modern Atomic TheoryChapter 10

  2. Rutherford’s Atom The concept of a nuclear atom (charged electrons moving around the nucleus) resulted from Ernest Rutherford’s experiments. • 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 elements arranged and how do they move?

  3. 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 • Electromagnetic radiation given off by atoms when they have been excited by any form of energy • flame tests • All waves are characterized by their velocity, wavelength, amplitude, and the number of waves that pass a point in a given time

  4. Electromagnetic Waves • velocity = c = speed of light • 2.997925 x 108 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 

  5. Types of Electromagnetic Radiation • Radiowaves =  > 0.01 m, low frequency and energy • Microwaves = 10-4m <  < 10-2m • Infrared (IR) • far = 10-4 <  < 10-5m • middle = 10-5 <  < 2 x 10-6m • near = 2 x 10-6 <  < 8 x 10-7m • Visible = 8 x 10-7 <  < 4 x 10-7m • ROYGBIV • Ultraviolet (UV) • near = 4 x 10-7 <  < 2 x 10-7m • far = 2 x 10-7 <  < 1 x 10-8m • X-rays = 10-8 <  < 10-10m • Gamma rays =  < 10-10

  6. Figure 10.4: The different wavelengths of electromagnetic radiation.

  7. Planck’s Revelation • Showed that light energy could be thought of as particles for certain applications • Stated that light came in particles called quanta or photons • Particles of light have fixed amounts of energy • Basis of quantum theory • The energy of the photon is directly proportional to the frequency of light • Higher frequency = More energy in photons

  8. 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!!

  9. 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

  10. 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

  11. Bohr’s Model • 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

  12. Bohr’s Model • Atoms have a minimum energy called the ground state • therefore they do not crash into the nucleus • The ground state of hydrogen corresponds to having its one electron in an energy level that is closest to the nucleus • Energy levels higher than the ground state are called excited states • 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

  13. 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

  14. 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

  15. 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 • 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

  16. 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

  17. Orbitals and Energy Levels • Principal energy levels identify how much energy the electrons in the orbital have • n • higher values mean orbital has higher energy • higher values mean orbital has farther average distance from the nucleus • 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

  18. Figure 10.21: The first four principal energy levels in the hydrogen atom.

  19. The relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

  20. The three 2p orbitals: (a) 2px, (b) 2pz, (c) 2py.

  21. The shapes and labels of the five 3d orbitals.

  22. 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

  23. Orbitals, Sublevels & Electrons • 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

  24. The Order of Filling of Orbitals

  25. 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 1s22s22p6= [Ne]

  26. Electron Configuration • 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

  27. Figure 10.29: The electron configurations in the sublevel last occupied for the first eighteen elements.

  28. Figure 10.30: Partial electron configurations for the elements potassium through krypton.

  29. s1 s2 p1 p2 p3 p4 p5 s2 1 2 3 4 5 6 7 p6 d1 d2d3 d4 d5 d6 d7 d8 d9 d10 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

  30. Figure 10.33: The positions of the elements considered in Example 10.3.

  31. The periodic table with atomic symbols, atomic numbers, and partial electron configurations.

  32. The periodic table with atomic symbols, atomic numbers, and partial electron configurations (cont’d).

  33. 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

  34. The classification of elements as metals, nonmetals, and metalloids.

  35. Metallic Character • Metalloids • Also known as semi-metals • Show some metal and some nonmetal properties • 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 • Metals • malleable & ductile • shiny, lustrous • conduct heat and electricity • most oxides basic and ionic • form cations in solution • lose electrons in reactions - oxidized

  36. 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

  37. 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

  38. Trend in Ionization Energy • Minimum energy needed to remove a valence electron from an atom • gas state • The lower the ionization energy, the easier it is to remove the electron • meatls have low ionization energies • Ionization Energy decreases down the group • valence electron farther from nucleus • Ionization Energy increases across the period • left to right

  39. Trend in 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

  40. Relative atomic sizes for selected atoms.