1. Chapter 5 Electrons In Atoms

2. Topics to Be Covered  5.1 Light and Quantized Energy 136-145  5.2 Quantum Theory and the Atom 146-155  5.3 Electron Configuration 156-162

3. Section 5.1 Light and Quantized Energy

4. The Atom & Unanswered Questions  Early 1900s Discovered 3 subatomic particles Continued quest to understand atomic structure Rutherford’s model  Positive charge in nucleus  Fast moving electrons around that  No accounting for differences and similarities in chemical behavior

5. The Atom and Unanswered Questions  Example: Lithium, sodium, and potassium have similar chemical behaviors (explained more in next chapter)  Early 1900s Scientists began to unravel mystery Certain elements emitted visible light when heated in a flame Analysis revealed chemical behavior depends on arrangement of electrons

6. The Wave Nature of Light  Electromagnetic radiation A form of energy that exhibits wavelike behavior as it travels through space Visible light is a type of ER

7. Characteristics of Waves  All waves can be described by several characteristics 1. Wavelength 2. Frequency 3. Amplitude

8. Wavelength  Represented by lambda λ  Shortest distance between equivalent points on a continuous waves  Measure crest to crest or trough to trough  Usually expressed in m, cm, or nm

9. Frequency  Represented by nu ν  The number of waves that pass a given point per second  Given in the unit of hertz (Hz)  1 Hz = 1 wave per second

10. Amplitude  The wave’s height from the origin to a crest or from the origin to a trough  Wavelength and frequency do not affect amplitude

11. Speed  All electromagnetic waves in a vacuum travel at a speed of 3.00 x 10 8 m/s This includes visible light  The speed of light has its own symbol C  C= λν

12. Electromagnetic Spectrum  Also called the EM spectrum  Includes all forms of electromagnetic radiation  With the only differences in the types of radiation being their frequencies and wavelengths

13. Electromagnetic Spectrum  Figure 5.5

14. Problems  Page 140  Calculating Wavelength of an EM Wave

15. Particle Nature of Light  Needed to explain other properties of light Heated objects emit only certain frequencies of light at a given temperature Some metals emit electrons when light of a specific frequency shines on them

16. Quantum Concept  When objects are heated they emit glowing light  1900 Max Planck began searching for an explanation  Studied the light emitted by heated objects  Startling conclusion

17. Quantum Concept  Planck discovered: Matter can gain or lose energy only in small specific amounts These amounts are called quanta Quantum—is the minimum amount of energy that can be gained or lost by an atom

18. Example  Heating a cup of water Most people thought that you can add any amount of thermal energy to the water by regulating the power and duration of the microwaves In actuality, the temperature increases in infinitesimal steps as its molecules absorb quanta of energy, which appear to be a continuous manner

19. Quantum Concept  Planck proposed that energy emitted by hot objects was quantized  Planck further demonstrated mathematically that a relationship exists between energy of a quantum and a frequency

20. Energy of a Quantum  E quantum =hv  E quantum represents energy  h is Planck’s constant  v represents frequency

21. Planck’s Constant  Symbol = h  6.626 x 10 -34 J*s  J is the symbol for joule The SI unit of energy  The equation shows that the energy of radiation increases as the radiation’s frequency, v, increases.

22. Planck’s Theory  For given frequencies Matter can emit/absorb energy only in whole number multiples of hv  1hv, 2hv, 3hv, 4hv etc. Matter can have only certain amounts of energy  Quantities of energy between these values do not exist

23. The Photoelectric Effect  Photoelectric effect electrons, called photoelectrons are emitted from a metal’s surface when light of a certain frequency, or higher than a certain frequency shines on the surface

24. Light’s Dual Nature  Einstein proposed in 1905 that light has a dual nature  photon—a massless particle that carries a quantum of energy

25. Energy of a Photon  E photon =hv  E photon represents energy  h is Planck’s constant  v represents frequency

26. Light’s Dual Nature  Einstein proposed Energy of a photon must have a certain threshold value to cause the ejection of a photoelectron from the surface of a metal Even small #s of photons with energy above the threshold value will cause the photoelectric effect  Einstein won Nobel Prize in Physics in 1921

27. Sample Problems  Page 143  Sample Problem 5.2  Calculating Energy of a Photon

28. Atomic Emission Spectrum Fe

29. Section 5.2 Quantum Theory and The Atom

30. Bohr’s Model of the Atom  Dual-nature explains more Atomic Emission Spectra Not continuous Only certain frequencies of light Explained the Atomic Emission Spectra

31. Energy States of Hydrogen  Bohr proposed certain allowable energy states  Bohr proposed electrons could travel in certain orbitals Fixed circular orbits ~ merry-go-round

32. Energy states of Hydrogen  Ground State Lowest allowable energy state of an atom  Orbit size Smaller the orbit, the lower the energy state/level Larger the orbit, the higher the energy state/level

33. Energy states of Hydrogen  Hydrogen can have many excited states It only has one electron  Quantum Number Number assigned to each orbital ~ level n

34. The Hydrogen Line Spectrum  Hydrogen Ground State Electron is in n=1 orbit Does not radiate energy  Hydrogen Excited State Energy is added to the atom from outside source Electron moves to a higher energy orbit

35. The Hydrogen Line Spectrum  Only certain atomic energies are possible Only certain frequencies of electromagnetic radiation can be emitted Frequencies in H are not evenly spaced Electrons are excited different amounts

36. The limits of Bohr’s Model  Bohr’s model explained H  Fails to explain any other element’s spectrum  Laid groundwork for atomic models to come  Fundamentally incorrect  Movements of e - are not completely understood even now

37. The Quantum Mechanical Model of the Atom  Mid 1920s Scientists convinced Bohr’s model is incorrect  De Broglie Electrons as waves

38. The Quantum Mechanical Model of the Atom  Electrons as waves Light possesses a dual nature De Broglie began to think of electrons as wave/particles

39. The Quantum Mechanical Model of the Atom  De Broglie equation Predicts that all moving particles have wave characteristics Explains impossibility for people to detect wavelengths of fast moving cars Electron having wavelike motion and restriction to a circular orbit of a fixed radius only certain wavelengths and frequencies are possible

40. The Heisenberg Uncertainty Principle  Heisenberg German theoretical physicist (1901-1976) Showed that it is impossible to take the measurement of an object without disturbing the object

41. The Heisenberg Uncertainty Principle  Heisenberg German theoretical physicist (1901-1976) Heisenberg Uncertainty Principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time  Cannot assigned fixed paths to electrons Only know a probability for an electron to occupy a certain region

42. The Schrödinger Wave Equation  Erwin Schrödinger furthered the wave- particle theory of De Broglie  Derived an equation that treated the H electron as a wave  Succeeded where Bohr’s failed  Quantum Mechanical Model of the Atom is the atomic model in which electrons are treated as waves AKA thewave mechanical model

43. Orbitals

44. Electron’s probable location  The wave function (Schrödinger) predicts a 3D region around the nucleus  This describes the probable location of the electron  This area is called the atomic orbital

45. Hydrogen’s Atomic Orbitals  Orbitals do not have a defined size Boundaries are fuzzy  There are four quantum numbers in atomic orbitals

46. Hydrogen’s Atomic Orbitals  Principal Quantum Number Principal quantum number (n) indicates the relative size and energy of atomic orbitals  As n increases, so does the size and energy  N specifies atoms major energy levels

47. Hydrogen’s Atomic Orbitals  Principal Quantum Number Principal energy level is each major energy level There are up to 7 energy levels  Energy Sublevels Principal energy levels contain energy sublevels N =1 has 1 sublevel N =2 has 2 sublevels etc.

48. Hydrogen’s Atomic Orbitals  Shapes of Orbitals Four types we will discuss SS PP DD FF

49. Hydrogen’s Atomic Orbitals  Shapes of Orbitals S orbital  Spherical in shape  Can contain 2 electrons P orbital  Dumbbell shape Px, Py, Pz  Can hold up to 6 electrons

50. Atomic Orbitals  Shape of Orbitals D orbital  Four look similar, X Fifth looks like a dumbbell with a ring  Can hold up to 10 electrons F orbital  Can hold up to 14 electrons

52. Section 5.3 Electron Configuration

53. Ground-State Electron Configuration  Electron configuration The arrangement of electrons in an atom Low-energy systems are more stable than high-energy systems Electrons assume arrangements that give lowest possible energy (ground state)

54. Ground-State Electron Configuration  There are three rules/principles that define how electrons can be arranged Aufbau Principle Pauli Exclusion Principle Hund’s Rule

55. The Aufbau Principle  Aufbau states that each electron occupies the lowest energy orbit available Sits closest to the nucleus possible

56. Pauli Exclusion Principle  Electrons in orbits can be represented by arrows. Electrons have either an upspin or a downspin.  Pauli Exclusion Principle states that a maximum of two electrons can occupy a single atomic orbital, but only if the electrons have opposite spins.

57. Hund’s Rule  Hund’s Rule states that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals.

58. Exceptions to predicted configurations  Chromium  Copper

59. Valence Electrons  Valence electrons are defined as electrons in the atom’s outermost orbital (generally the highest principal energy level) Tend to be in s and p orbitals for the the first three energy levels  Electron-dot structure consists of the element’s symbol which represents the nucleus and inner energy levels

60. Valence Electrons  Electron-dot structure consists of the element’s symbol which represents the nucleus and inner energy levels Surrounded by dots that represent the valence electrons