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Introductory Chemistry , 2 nd Edition Nivaldo Tro

Introductory Chemistry , 2 nd Edition Nivaldo Tro. Chapter 9 Electrons in Atoms and the Periodic Table. Why do Blimps Float?. Because they are filled with a gas less dense than air Early blimps used hydrogen gas; hydrogen’s flammability led to the Hindenburg disaster

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Introductory Chemistry , 2 nd Edition Nivaldo Tro

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  1. Introductory Chemistry, 2nd EditionNivaldo Tro Chapter 9 Electrons in Atoms and the Periodic Table

  2. Why do Blimps Float? • Because they are filled with a gas less dense than air • Early blimps used hydrogen gas; hydrogen’s flammability led to the Hindenburg disaster • Blimps now use helium, a nonflammable gas – in fact it doesn’t undergo any chemical reactions • This chapter investigates models of the atom we use to explain the differences in the properties of the elements Tro's Introductory Chemistry, Chapter 9

  3. Electromagnetic Radiation • Light is one of the forms of energy • Light is one type of a more general form of energy called electromagnetic radiation • Electromagnetic radiation travels in waves Tro's Introductory Chemistry, Chapter 9

  4. Characteristics of a Wave • Wavelength = distance from peak to peak • Amplitude = height of the peak • Frequency = the number of wave peaks that pass in a given time • Speed = rate the waves travel Tro's Introductory Chemistry, Chapter 9

  5. Particles of Light • Scientists in the early 20th century showed that electromagnetic radiation was composed of particles we call photons • Max Planck and Albert Einstein • photons are particles of light energy • Each wavelength of light has photons that have a different amount of energy • the longer the wavelength, the lower the energy of the photons Tro's Introductory Chemistry, Chapter 9

  6. The Electromagnetic Spectrum • Light passed through a prism is separated into all its colors = continuous spectrum; colors blend into each other • Color of light is determined by its wavelength Tro's Introductory Chemistry, Chapter 9

  7. Electromagnetic Spectrum Visible light is a very small portion of the electromagnetic spectrum Tro's Introductory Chemistry, Chapter 9

  8. Light’s Relationship to Matter • Atoms can absorb energy, but they must eventually release it • When atoms emit energy, it is released in the form of light = emission spectrum • Atoms don’t absorb or emit all colors, only very specific wavelengths; the spectrum of wavelengths can be used to identify the element Tro's Introductory Chemistry, Chapter 9

  9. Emission Spectrum or Line Spectrum Tro's Introductory Chemistry, Chapter 9

  10. Line Spectra = specific wavelengths are emitted; characteristic of atoms Tro's Introductory Chemistry, Chapter 9

  11. The Bohr Model of the Atom • Nuclear Model of atom does not explain how atom can gain or lose energy • Neils Bohr developed a model to explain how structure of the atom changes when it undergoes energy transitions • Bohr postulated that energy of the atom was quantized, and that the amount of energy in the atom was related to the electron’s position in the atom • quantized means that the atom could only have very specific amounts of energy Tro's Introductory Chemistry, Chapter 9

  12. Bohr Model of Atom: Electron Orbits • In the Bohr Model, electrons travel in orbits or energy levels around the nucleus • The farther the electron is from the nucleus the more energy it has

  13. The Bohr Model of the Atom:Orbits and Energy • Each orbit (energy level) has a specific amount of energy • Energy of each orbit is symbolized by n, with values of 1, 2, 3 etc; the higher the value the farther it is from the nucleus and the more energy an electron in that orbit has Tro's Introductory Chemistry, Chapter 9

  14. The Bohr Model of the Atom:Energy Transitions • Electrons can move from a lower to a higher (farther from nucleus) energy level by absorbing energy • When the electron moves from a higher to a lower (closer to nucleus) energy level, energy is emitted from the atom as a photon of light Tro's Introductory Chemistry, Chapter 9

  15. The Bohr Model of the AtomGround and Excited States • Ground state – atoms with their electrons in the lowest energy level possible; this lowest energy state is the most stable. • Excited state – a higher energy state; electrons jump to higher energy levels by absorbing energy • Atom is less stable in an excited state; it will release the extra energy to return to the ground state Tro's Introductory Chemistry, Chapter 9

  16. Electron Energy Levels: Energy LevelHow many e fit? (2n2) 3rd 18 electrons 2 x 32 2nd 8 electrons 2 x 22 1st 2 electrons 2 x 12 Each energy level has a maximum # of electrons it can hold. H has one electron; it is in the 1st energy level. Bohr model H Tro's Introductory Chemistry, Chapter 9

  17. Bohr Model for AtomElectrons fill the Lowest energy levels first C Bohr Model for C with 6 electrons Tro's Introductory Chemistry, Chapter 9

  18. The Bohr Model of the AtomSuccess and Failure • The Bohr Model very accurately predicts the spectrum of hydrogen with its one electron • It is inadequate when applied to atoms with many electrons • A better theory was needed Tro's Introductory Chemistry, Chapter 9

  19. The Quantum-Mechanical ModelOrbitals • Erwin Schrödinger used mathematics to predict probability of finding an electron at a certain location in the atom • Result is a map of regions in the atom that have a particular probability for finding the electron • Orbital = a region with a very high probability of finding the electron when it has a particular amount of energy Tro's Introductory Chemistry, Chapter 9

  20. The Quantum-Mechanical Model • Each principal energy level or shell has one or more subshells • # of subshells same as the principal quantum number or shell • The subshells are often represented as a letter • s, p, d, f • Each kind of subshell has orbitals with a particular shape Tro's Introductory Chemistry, Chapter 9

  21. Shells & Subshells Tro's Introductory Chemistry, Chapter 9

  22. Probability Maps & Orbital Shapesorbitalsare spherical Tro's Introductory Chemistry, Chapter 9

  23. Probability Maps & Orbital Shapeporbitals Tro's Introductory Chemistry, Chapter 9

  24. Subshells and Orbitals • The subshells of a principal shell have slightly different energies • the subshells in a shell of H all have the same energy, but for multielectron atoms the subshells have different energies • s < p < d < f • Each subshell contains one or more orbitals • s subshells have 1 orbital • p subshells have 3 orbitals • d subshells have 5 orbitals • f subshells have 7 orbitals Tro's Introductory Chemistry, Chapter 9

  25. The Quantum Mechanical ModelEnergy Transitions • As in Bohr Model, atoms gain or lose energy as electron moves between orbitals in different energy shells and subshells • The ground state of the electron is the lowest energy orbital it can occupy • Excited state = when an electron moves to a higher energy orbital Tro's Introductory Chemistry, Chapter 9

  26. The Bohr Model vs.The Quantum Mechanical Model • Both the Bohr and Quantum Mechanical models predict the spectrum of hydrogen very accurately • Only the Quantum Mechanical model predicts the spectra of multielectron atoms Tro's Introductory Chemistry, Chapter 9

  27. Electron Configurations • Electron configuration = distribution of electrons into the various energy shells and subshells in an atom in its ground state • Each energy shell and subshell has a maximum number of electrons it can hold • s = 2, p = 6, d = 10, f = 14 Tro's Introductory Chemistry, Chapter 9

  28. Writing Electron Configurations • We place electrons in the energy shells and orbitals in order of energy, from low energy up: Aufbau Principle (order of filling of orbitals) • The d and f orbitals overlap into the higher energy levels Tro's Introductory Chemistry, Chapter 9

  29. 6d 7s 5f 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s Energy Relative Energy of Orbitals in the Quantum Mechanical Model Tro's Introductory Chemistry, Chapter 9

  30. Order of Subshell Fillingin Ground State Electron Configurations 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) 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s next, draw arrows through the diagonals, looping back to the next diagonal each time Tro's Introductory Chemistry, Chapter 9

  31. Filling the Orbitals in a Subshellwith Electrons • Energy shells fill from lowest energy to high • Subshells fill from lowest energy to high • s → p → d → f • A single orbital can hold a maximum of 2 electrons (Pauli’s exclusion principle); orbitals that are in the same subshell have the same energy • When filling orbitals that have the same energy, place one electron in each before completing pairs (Hund’s rule) Tro's Introductory Chemistry, Chapter 9

  32. Electron Configuration of Atoms in their Ground State • Electron configuration = order of filling with electrons; number of electrons in that subshell written as a superscript Kr = 36 electrons= 1s22s22p63s23p64s23d104p6 • Shorthand way: use the symbol of the previous noble gas in brackets to represent all the inner electrons, then just write the last set Rb = 37 electrons= 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1 Tro's Introductory Chemistry, Chapter 9

  33. Electron Configurations how many electrons in that orbital • Nitrogen: 1s22s22p3 energy level orbital (atomic number = 7) Tro's Introductory Chemistry, Chapter 9

  34. Example – Write the Ground State Orbital Diagram and Electron Configuration of Magnesium. • Determine the atomic number of the element from the Periodic Table • This gives the number of protons and electrons in the atom Mg, Z = 12, so Mg has 12 protons and 12 electrons Tro's Introductory Chemistry, Chapter 9

  35. 1s 2s 2p 3s 3p Example – Write the Ground State Orbital Diagram and Electron Configuration of Magnesium. • Draw 9 boxes to represent the first 3 energy levels sandp orbitals Tro's Introductory Chemistry, Chapter 9

  36. Example – Write the Ground State Orbital Diagram and Electron Configuration of Magnesium. • Add one electron to each box in a set, then pair the electrons before going to the next set until you use all the electrons • When pairing, put in opposite arrows          1s 2s 2p 3s 3p Tro's Introductory Chemistry, Chapter 9

  37.  1s  2s  2p    3s 3p Example – Write the Ground State Orbital Diagram and Electron Configuration of Magnesium. • Use the diagram to write the electron configuration • Write the number of electrons in each set as a superscript next to the name of the orbital set 1s22s22p63s2 = [Ne]3s2 Tro's Introductory Chemistry, Chapter 9

  38. Valence Electrons • Valence electrons = electrons in all the subshells with the highest principal energy shell (outermost shell) • Core electrons = in lower energy shells • Valence electrons responsible for both chemical and physical properties of atoms. • Valence electrons responsible for chemical reactions Tro's Introductory Chemistry, Chapter 9

  39. Valence Electrons Rb = 37 electrons= 1s22s22p63s23p64s23d104p65s1 • The highest principal energy shell of Rb that contains electrons is the 5th, therefore Rb has 1 valence electron and 36 core electrons Kr = 36 electrons= 1s22s22p63s23p64s23d104p6 • The highest principal energy shell of Kr that contains electrons is the 4th, therefore Kr has 8 valence electrons and 28 core electrons Tro's Introductory Chemistry, Chapter 9

  40. How many valence electrons does each atom have? carbon: 1s22s22p2 chlorine: 1s22s22p63s23p5 Tro's Introductory Chemistry, Chapter 9

  41. How many valence electrons does each atom have? carbon: 1s22s22p2= 4 chlorine: 1s22s22p63s23p5= 7 Tro's Introductory Chemistry, Chapter 9

  42. Electron Configurations andthe Periodic Table Tro's Introductory Chemistry, Chapter 9

  43. Electron Configurations fromthe Periodic Table • Elements in the same period (row) have valence electrons in the same principal energy shell • The number of valence electrons increases by one as you progress across the period • Elements in the same group (column) have the same number of valence electrons and they are in the same kind of subshell Tro's Introductory Chemistry, Chapter 9

  44. Electron Configuration & the Periodic Table • Elements in the same column have similar chemical and physical properties because their valence shell electron configuration is the same • The number of valence electrons for the main group elements is the same as the group number Tro's Introductory Chemistry, Chapter 9

  45. The Explanatory Power ofthe Quantum-Mechanical Model • The properties of the elements are largely determined by the number of valence electrons they contain • Since elements in the same column have the same number of valence electrons, they show similar properties Tro's Introductory Chemistry, Chapter 9

  46. The Noble Gas Electron Configuration • The noble gases have 8 valence electrons • except for He, which has only 2 electrons • Noble gases are especially unreactive • He and Ne are practically inert • Reason noble gases are unreactive is that the electron configuration of the noble gases is especially stable Tro's Introductory Chemistry, Chapter 9

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