Unraveling Atomic Structure: Light and Energy Waves

1. Chapter 12 Electrons in Atoms

2. Although three subatomic particles had been discovered by the early 1900s, the quest to understand the atom and its structure had just begun. Rutherford proposed that all of an atom’s positive charge and virtually all of its mass are concentrated in a nucleus that is surrounded by fast-moving electrons.

3. Although his nuclear model was a major scientific development, it lacked detail about how electrons occupy the space surrounding the nucleus. ? ? ? ? ? ? ? ? ? ? ? Because of this, scientists in the early twentieth century found Rutherford’s nuclear atomic model to be fundamentally incomplete.

4. Nor did it address the question of why the negatively charged electrons are not pulled into the atom’s positively charged nucleus. - - - - - - -

5. Chemists found Rutherford’s nuclear model lacking because it did not begin to account for the differences in chemical behavior among the various elements. Alkali Metal Family

6. In the early 1900s, scientists began to unravel the puzzle of chemical behavior. They had observed that certain elements emitted visible light when heated in a flame. What caused these differences?

7. ELECTRONS! Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms. When excited electrons drop to lower energy levels they release light!

8. In order for you to better understand this relationship and the nature of atomic structure, it will be helpful for you to first understand the nature of light.

10. Light and Quantized Energy • Wave Nature of light • Visible light is a form of energy that exhibits wave-like properties known as electromagnetic radiation. • Electromagnetic radiation – A series of electromagnetic waves that travel in a vacuum at a speed of 3.0 x 108 m/s. • Radio waves • Microwaves • Infrared waves • Visible light • Ultraviolet rays • X-rays • Gamma rays

11. Electromagnetic Spectrum

12. The Electromagnetic Spectrum Radio Waves- Radio, TV, Cell phones, Police Radio Microwaves Doppler Radar, Microwaves, some cell phones Infrared Waves Heat, thermal imaging, alarm systems, remote controls-TV

13. ROY G. BIV • ROY G. BIV is an acronym that helps us remember the order of the visible light spectrum! • Red, orange, yellow, green, blue, indigo, violet. • As you approach violet, the frequency (and energy) of the wave increases

14. Electromagnetic Spectrum Ultraviolet Waves UV lamps- sterilization Tanning Some insects can see UV radiation X rays X rays Airport secutrity Gamma Rays Radiation Treatment Medical Tracers

15. Light and Quantized Energy • What are waves • Mechanical waves – require a medium to travel (air, water, or rope). • Electromagnetic waves – no medium • Matter waves – particles and electrons

16. Light and Quantized Energy • Wave Properties • Wavelength (λ or lambda) – distance between equivalent points • 1nm = 1 × 10-9 m • Amplitude – height from origin to crest, involves the intensity of the light

17. Light and Quantized Energy • Wave Properties • Frequency ( or nu) – how many waves pass a given point per second. • 1 hertz (Hz) = 1 wave per second • 1MHz = 1  106 Hz

18. Light and Quantized Energy • Wave Nature of light • All electromagnetic light moves at the speed of 3.00 × 108 m/s and is represented by the symbol, c. • The speed of light is the product of the wavelength (λ) and frequency ().

19. Light and Quantized Energy • Although the speed of electromagnetic waves are constant, the frequency and the wavelength may vary. • As you can see from the equation, wavelength and frequency are inversely related; in other words, as one quantity increases, the other decreases.

20. Light and Quantized Energy

21. Light and Quantized Energy • What is the wavelength of a microwave having a frequency of 3.44 x 109 Hz? = 8.72 × 10-2 m

22. Light and Quantized Energy • A helium-neon laser emits light with a wavelength of 633 nm. What is the frequency of this light? c = 3.00  108 m/s 1 nm = 1  10-9 m λ = 633 nm 633 nm = 6.33  10-7 m

23. Light and Quantized Energy • Particle Nature of Light • While considering light as a wave does explain much of its everyday behavior, it fails to adequately describe important aspects of light’s interactions with matter. • Glowing substances • Photoelectric effect

24. Light and Quantized Energy • Particle Nature of Light • The wave model of light cannot explain why heated objects emit only certain frequencies of light at a given temperature, or why some metals emit electrons when colored light of a specific frequency shines on them. • In other words, how is the energy of a photon determined by its frequency (color)?

25. 1858–1947 Light and Quantized Energy • Particle Nature of Light • In 1900, the German physicist Max Planck began searching for an explanation as he studied the light emitted from heated objects. • matter can gain or lose energy only in small, specific amounts called quanta. • Quantum – minimum amount of energy that can be gained or lost by an atom. • A quantum of light is known as a photon This is light acting like a particle!

26. Light and Quantized Energy • Particle Nature of Light • Planck found that the energy of a quantum of energy (photon) is directly proportional to the frequency. h = 6.626 × 10-34 J·s

27. Light and Quantized Energy • What is the energy of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 1014 Hz? E = (6.626 × 10-34 J·s)(7.23 × 1014 s-1) E = 4.79 × 10-19 J

28. Light and Quantized Energy • Particle Nature of Light- Einstein • Photoelectric effect – electrons, called photoelectrons, will be emitted from a metal when light above a certain frequency is shined on it.

29. Light and Quantized Energy • Einstein's explanation treated light like a particle. Unless the incoming light has a high enough frequency (energy) it can’t release the photoelectron.

30. Light and Quantized Energy • Particle Nature of Light • Atomic emission spectrum – non-continuous spectra emitted by glowing atoms. • Conclusion: the different lines of the spectra are explained by the different energies between orbitals

31. How did we get to this??? Light seems most obviously a wave The major problem is that there seems to be no medium. Light can travel through the vacuum of space, in this sense light is clearly like a particle. Atoms can only absorb energy in discrete units, many small bursts of energy that add up to the same thing will not do the trick. Light must then come a discrete units, like particles. The name given to a quantum of light is the photon. The energy of a photon is determined by its frequency (color)

32. Quantum Theory and the Atom • Energy levels – electrons orbit in circles around the nucleus at fixed energy amounts (quantized). • Ground State – atoms at the lowest energy levels • The higher the energy level the farther it is from the nucleus. Quantum

33. Quantum Theory and the Atom • When an atom gains energy, it is said to be in an excited state. • And although a hydrogen atom contains only a single electron, it is capable of having many different excited states.

34. Quantum Theory and the Atom • Building on Planck’s and Einstein’s concepts of quantized energy (quantized means that only certain values are allowed), Bohr proposed that the hydrogen atom has only certain allowable energy states. • Impressively, Bohr’s model also correctly predicted the frequencies of the lines in hydrogen’s atomic emission spectrum.

35. Let’s Compare… • Let’s compare this to a ladder. • The lowest rung on the ladder represents the lowest energy level. • A person can climb up or down the ladder by going from rung to rung • A person on a ladder cannot stand in between the rungs. • To move from one energy level to another, an electron must gain or lose just the right amount of energy just like on a ladder a person must move just the right distance

36. Quantum Theory and the Atom

37. Quantum Theory and the Atom Bohr’s model worked well for hydrogen, but… Fell apart for every other atom on the periodic table!!! It did, however, point in the right direction!

38. Quantum Theory • REVIEW- A Quantum of energy is the amount of energy required to move an electron from its present energy level to the next higher one. • The higher the energy level, the farther from the nucleus it is. • The amount of energy gained or lost by an electron is not always the same. Unlike the rungs of a ladder, the energy levels in an atom are not equally spaced. • The energy levels become more closely spaced the greater the distance from the nucleus • Escalator

39. 1892–1987 Quantum Theory and the Atom • In 1924, a young French graduate student in physics named Louis de Broglie proposed an idea that eventually accounted for the fixed energy levels of Bohr’s model. • If waves could be treated like a particle, could particles be treated like waves?

40. Louis de Broglie • Louis de Broglie (1924) suggested that electrons are also waves (not particles) • This can be difficult to comprehend: normally we perceive objects as solid. • The reason objects seem solid is because they have a small wave length. small mass large mass

41. = =  • De Broglie knew that if an electron has wavelike motion and is restricted to circular orbits of fixed radius, the electron is allowed only certain possible wavelengths, frequencies, and energies. In other words, it would be quantized just like observed.

42. Quantum Theory and the Atom • Developing his idea, de Broglie derived an equation for the wavelength (λ) of a particle of mass (m) moving at velocity (ν). Experiments show that the smaller the particle, the more important it’s wave properties!

43. 1901–1976 Quantum Theory and the Atom • Step by step, scientists such as Rutherford, Bohr, and de Broglie had been unraveling the mysteries of the atom. • However, a conclusion reached by the German theoretical physicist Werner Heisenberg a contemporary of de Broglie, proved to have profound implications for atomic models.

44. Quantum Theory and the Atom • Heisenberg Uncertainty Principle • You can’t precisely know both the position and velocity of a particle at the same time. No, but I know where I’m at! Do you know how fast you were going? When Heisenberg was pulled over for speeding

45. 1887–1961 Quantum Theory and the Atom • In 1926, Austrian physicist Erwin Schrödinger furthered the wave-particle theory proposed by de Broglie. Schrödinger derived an equation that treated the hydrogen atom’s electron as a wave

46. Quantum Theory and the Atom • Remarkably, unlike Bohr’s model, Schrödinger’s new model for the hydrogen atom seemed to apply equally well to all elements! With this equation, the modern Quantum Mechanical Model was born.

47. Quantum Mechanical Model Quantum Theory and the Atom • Electron position and energy are described using energy levels, energy sublevels, orbital shapes, and spin. nucleus electron cloud 90% probability of finding the electron within this space

48. Quantum Mechanical Model

49. Quantum Theory and the Atom • Principal Energy Level (n) • Describes distance from the nucleus and general energy. • n = 1, 2, 3, 4, …. • The higher the energy level the greater the average distance from the nucleus. • Each energy level contains sublevels • The number of sublevels on a level is equal to the energy level (n). • 1st energy level has 1 sublevel • 2nd energy level has 2 sublevels

50. Quantum Theory and the Atom • Each sublevel contains orbitals. • orbital: a three-dimensional region around the nucleus in which an electron moves and is found 90% of the time. • Each orbital can hold up to two electrons. • The total number of orbitals on a level = n2. • Each sublevel has a different shape of orbital on the level. • These shapes are represented by the symbols s, p, d, or f.