2. Section 5-1 Section 5.1 Light and Quantized Energy Compare the wave and particle natures of light. radiation: the rays and particles —alpha particles, beta particles, and gamma rays—that are emitted by radioactive material Define a quantum of energy, and explain how it is related to an energy change of matter. Contrast continuous electromagnetic spectra and atomic emission spectra.

3. Section 5-1 Section 5.1 Light and Quantized Energy (cont.) electromagnetic radiation wavelength frequency amplitude electromagnetic spectrum Light, a form of electronic radiation, has characteristics of both a wave and a particle. quantum Planck's constant photoelectric effect photon atomic emission spectrum

4. Section 5-1 The Atom and Unanswered Questions Recall that in Rutherford's model, the atom’s mass is concentrated in the nucleus and electrons move around it. The model doesn’t explain how the electrons were arranged around the nucleus. The model doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus.

5. Section 5-1 The Atom and Unanswered Questions (cont.) In the early 1900s, scientists observed certain elements emitted visible light when heated in a flame. Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms.

6. Section 5-1 The Wave Nature of Light Visible light is a type of electromagnetic radiation, a form of energy that exhibits wave-like behavior as it travels through space.electromagnetic radiation All waves can be described by several characteristics.

7. Section 5-1 The Wave Nature of Light (cont.) The wavelength (λ) is the shortest distance between equivalent points on a continuous wave.wavelength The frequency (f) is the number of waves that pass a given point per second. The unit for frequency is 1/sec or sec-1, which is known as a Hertz.frequency The amplitude is the wave’s height from the origin to a crest.amplitude

8. Section 5-1 The Wave Nature of Light (cont.)

9. Section 5-1 The Wave Nature of Light (cont.) The speed of light (3.00  10 8 m/s) is the product of it’s wavelength and frequency c = λf. c = f

10. EM Spectrum GIVEN: f = ? = 434 nm = 4.34  10 -7 m c = 3.00  10 8 m/s WORK : f = c f = 3.00  10 8 m/s 4.34  10 -7 m f = 6.91  10 14 Hz EX: Find the frequency of a photon with a wavelength of 434 nm.

11. Section 5-1 The Wave Nature of Light (cont.) Sunlight contains a continuous range of wavelengths and frequencies. A prism separates sunlight into a continuous spectrum of colors. The electromagnetic spectrum includes all forms of electromagnetic radiation.electromagnetic spectrum

12. Section 5-1 The Wave Nature of Light (cont.)

13. Section 5-1 The Particle Nature of Light The wave model of light cannot explain all of light’s characteristics. Matter can gain or lose energy only in small, specific amounts called quanta. Max Planck (1900) Observed - emission of light from hot objects Concluded - energy is emitted in small, specific amounts (quanta) A quantum is the minimum amount of energy that can be gained or lost by an atom.quantum Planck’s constant has a value of 6.626  10 –34 J ● s.Planck’s constant

14. Section 5-1 The Particle Nature of Light (cont.) The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.photoelectric effect

15. Section 5-1 The Particle Nature of Light (cont.) Albert Einstein proposed in 1905 that light has a dual nature. A beam of light has wavelike and particle like properties. A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.photon E photon = hfE photon represents energy. h is Planck's constant. f represents frequency. E = h f

16. Quantum Theory GIVEN: E = ? f = 4.57  10 14 Hz h = 6.6262  10 -34 J·s WORK : E = hf E = ( 6.6262  10 -34 J·s )( 4.57  10 14 Hz ) E = 3.03  10 -19 J EX: Find the energy of a red photon with a frequency of 4.57  10 14 Hz.

17. Section 5-1 Atomic Emission Spectra Light in a neon sign is produced when electricity is passed through a tube filled with neon gas and excites the neon atoms. The excited atoms emit light to release energy.

18. Section 5-1 Atomic Emission Spectra (cont.)

19. Section 5-1 Atomic Emission Spectra (cont.) The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element.atomic emission spectrum Each element’s atomic emission spectrum is unique. Like a finger print.

20. Section 5-2 Section 5.2 Quantum Theory and the Atom (cont.) ground state quantum number de Broglie equation Heisenberg uncertainty principle Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals. quantum mechanical model of the atom atomic orbital principal quantum number principal energy level energy sublevel

21. Review of 5.1: All waves have distinct amplitudes, frequency, and wavelengths. All electromagnetic waves travel at the speed of light. C = (3.0x10 8 m/s) C = f The relationship between the energy and frequency is E = hf All elements have their own individual atomic emission spectrum and absorption spectrum.

22. Section 5-2 Bohr's Model of the Atom The dual wave-particle model of light could not explain all phenomena of light. In particular, it couldn’t explain why emission spectra was discontinuous. Niels Bohr, a Danish physicist that worked in Rutherford’s lab, proposed a model that would match hydrogen’s emission spectrum.

23. In Bohr’s model, the atom has only certain allowable energy states. (circular orbits) The lowest allowable energy state of an atom is called its ground state.ground state When an atom gains energy, it is in an excited state.

24. Section 5-2 Bohr's Model of the Atom (cont.) In Bohr’s model, atom’s with smaller energy states will have electrons in smaller orbits. Even Hydrogen can have many different excited states, depending on where its one electron is located.

25. Section 5-2 Bohr's Model of the Atom (cont.) Each orbit was given a number, n, called the principal quantum number.quantum number

26. Section 5-2 Bohr's Model of the Atom (cont.) Hydrogen’s single electron is in the n = 1 orbit in the ground state. No energy radiates in this state. When energy is added, the electron moves to the n = 2 orbit. This raises the electron to an excited state. If the electron would then drop from the higher orbit to a lower one, it will release energy in the form of a photon as it drops. The photon corresponds to the energy difference between the two levels.

27. (p.148) Because there are only certain energy levels, there are only certain frequencies of radiation that can be emitted. E = hf Think of this like rungs of a ladder. Just as you can only go up or down from rung to rung, an electron can only move from one orbit to another. Unlike ladder rungs, however, the energy levels are not evenly spaced.

28. Electrons that drop from higher-energy orbits down to the 2 nd orbit make all hydrogen’s visible lines—the Balmer series. Electrons in Hydrogen that drop from higher levels to the 1 st orbit release ultraviolet light and are in the Lyman series. Electrons in Hydrogen that drop from higher orbits to the 3 rd orbit radiate infrared and are in the Paschen series.

29. Section 5-2 Bohr's Model of the Atom (cont.)

30. Section 5-2 Bohr's Model of the Atom (cont.)

31. Section 5-2 Bohr's Model of the Atom (cont.) Bohr’s model explained the hydrogen’s spectral lines, but failed to explain any other element’s lines. The behavior of electrons is still not fully understood, but it is known they do not move around the nucleus in circular orbits.

32. Section 5-2 The Quantum Mechanical Model of the Atom Mid 1920s, Frenchman Louis de Broglie (1892–1987) hypothesized that particles, including electrons, could also have wavelike behaviors. He compared light waves to waves made on musical instruments with fixed ends. On the instruments, only multiples of half wavelengths are possible. Similarly, de Broglie reasoned that only odd numbers of wavelengths are allowed in a circular orbit with a fixed radius.

33. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) The figure illustrates that electrons orbit the nucleus only in whole-number wavelengths.

34. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) *de Broglie proposed that if waves can act like particles, the reverse must also be true. The de Broglie equation predicts that all moving particles have wave characteristics.de Broglie equation represents wavelengths h is Planck's constant. m represents mass of the particle. represents velocity.

35. German Werner Heisenberg (1901-1976), showed it is impossible to take any measurement of an object without disturbing it. Heisenberg compared trying to measure an electron’s position to trying to find a helium-filled balloon in a darkened room. The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.Heisenberg uncertainty principle The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.

36. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) When a photon interacts with an electron at rest, both the velocity and position of the electron are modified.

37. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) 1926, Austrian Erwin Schrodinger (1887-1961) continued the wave- particle theory. Schrödinger treated electrons as waves in a model called the quantum mechanical model of the atom.quantum mechanical model of the atom Like Bohr’s model, this model limits an electron’s energy to certain values. It does not try to describe the electron’s path. His model is a very complex model using wave functions. Schrödinger’s equation applied equally well to elements other than hydrogen.

38. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) The wave function predicts a three-dimensional region around the nucleus called the atomic orbital.atomic orbital The density at a given point is proportional to the probability of finding an electron at that point.

39. Section 5-2 Hydrogen Atomic Orbitals Just as Bohr’s model had numbers assigned to electron orbits, so does the quantum mechanical model. There are 4 quantum numbers for orbitals. Principal quantum number (n) indicates the relative size and energy of atomic orbitals.Principal quantum number n specifies the atom’s major energy levels, called the principal energy levels. As n increases, the orbital is larger and the atom’s energy increases. The lowest principle energy level has a principal quantum number of 1. principal energy levels A hydrogen atom in ground state will have a single electron in the n = 1 orbital.

40. Section 5-2 Hydrogen Atomic Orbitals (cont.) Energy sublevels are contained within the principal energy levels.Energy sublevels The numbers of sublevels in a principal energy level increase as n increase, much like there are more seats per row as you go higher up in a stadium.

41. Shapes of orbitals: –All s orbitals are spherical –All p orbitals are dumbbell-shaped –d and f orbitals do not all have the same shape. –Orbitals in higher sublevels are bigger than ones in lower sublevels. –(2s is bigger than 1s)

42. Section 5-2 Hydrogen Atomic Orbitals (cont.) Each energy sublevel relates to orbitals of different shape.

46. Section 5-2 Hydrogen Atomic Orbitals (cont.) The number of orbitals related to each sublevel is always an odd number. The maximum number of orbitals for each principal energy level equals n 2.

47. Section 5-3 Section 5.3 Electron Configuration (cont.) electron configuration aufbau principle Pauli exclusion principle Hund's rule valence electrons electron-dot structure A set of three rules determines the arrangement in an atom.

48. Section 5-3 Ground-State Electron Configuration The arrangement of electrons in the atom is called the electron configuration.electron configuration The aufbau principle states that each electron occupies the lowest energy orbital available.aufbau principle

49. Section 5-3 Ground-State Electron Configuration (cont.)

50. Section 5-3 Ground-State Electron Configuration (cont.) The Pauli exclusion principle states that a maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins.Pauli exclusion principle 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 energy level orbitals.Hund’s rule

52. Section 5-3 Ground-State Electron Configuration (cont.)

53. Section 5-3 Ground-State Electron Configuration (cont.) Noble gas notation uses noble gas symbols in brackets to shorten inner electron configurations of other elements.

54. Shorthand Configuration –Core e - : Go up one row and over to the Noble Gas. –Valence e - : On the next row, fill in the # of e - in each sublevel. –Example - Germanium [Ar]4s 2 4p 2 3d 10

55. Full energy level (noble gases) Full sublevel (s, p, d, f) Half-full sublevel Stability

56. Electron Configuration Exceptions –Copper EXPECT :[Ar] 4s 2 3d 9 ACTUALLY :[Ar] 4s 1 3d 10 –Copper gains stability with a full d-sublevel. D. Stability

57. Electron Configuration Exceptions –Chromium EXPECT :[Ar] 4s 2 3d 4 ACTUALLY :[Ar] 4s 1 3d 5 –Chromium gains stability with a half-full d- sublevel. D. Stability

58. Section 5-3 Valence Electrons Valence electrons are defined as electrons in the atom’s outermost orbitals— those associated with the atom’s highest principal energy level. (i.e. the s and p sublevels)Valence electrons Electron-dot structure consists of the element’s symbol representing the nucleus, surrounded by dots representing the element’s valence electrons.Electron-dot structure

59. Section 5-3 Valence Electrons (cont.)