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Chapter 10

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  1. Chapter 10 Electrons in Atoms

  2. Democritus (400 B.C.) • Proposed that matter was composed of tiny indivisible particles • Not based on experimental data • Greek: atomos

  3. Alchemy (next 2000 years) • Mixture of science and mysticism. • Lab procedures were developed, but alchemists did not perform controlled experiments like true scientists.

  4. John Dalton (1807) • British Schoolteacher • based his theory on others’ experimental data • Billiard Ball Model • atom is a uniform, solid sphere

  5. Henri Becquerel (1896) • Discovered radioactivity • spontaneous emission of radiation from the nucleus • Three types: • alpha () - positive • beta () - negative • gamma () - neutral

  6. J. J. Thomson (1903) • Cathode Ray Tube Experiments • beam of negative particles • Discovered Electrons • negative particles within the atom • Plum-pudding Model

  7. J. J. Thomson (1903) Plum-pudding Model • positive sphere (pudding) with negative electrons (plums) dispersed throughout

  8. Ernest Rutherford (1911) • Gold Foil Experiment • Discovered the nucleus • dense, positive charge in the center of the atom • Nuclear Model

  9. Rutherford’s Gold Foil Experiment

  10. (a) The results that the metal foil experiment would have yielded if the plum pudding model had been correct. (b) Actual results.

  11. Ernest Rutherford (1911) • Nuclear Model • dense, positive nucleus surrounded by negative electrons

  12. Niels Bohr (1913) • Bright-Line Spectrum • tried to explain presence of specific colors in hydrogen’s spectrum • Energy Levels • electrons can only exist in specific energy states • Planetary Model

  13. Niels Bohr (1913) • Planetary Model • electrons move in circular orbits within specific energy levels Bright-line spectrum

  14. Erwin Schrödinger (1926) • Quantum mechanics • electrons can only exist in specified energy states • Electron cloud model • orbital: region around the nucleus where e- are likely to be found

  15. Erwin Schrödinger (1926) Electron Cloud Model (orbital) • dots represent probability of finding an e-not actual electrons

  16. James Chadwick (1932) • Discovered neutrons • neutral particles in the nucleus of an atom • Joliot-Curie Experiments • based his theory on their experimental evidence

  17. James Chadwick (1932) Neutron Model • revision of Rutherford’s Nuclear Model

  18. Electromagnetic Radiation • Electromagnetic radiation – radiowaves, X-rays, microwaves, infrared waves, visible light, ultraviolet waves and gamma rays. • All electromagnetic radiation travel at the speed of light (c = 3.0 x 108 m/s) in a vacuum.

  19. The different wavelengths of electromagnetic radiation.

  20. The electromagnetic spectrum.

  21. Physics and the Quantum Mechanical Model • Amplitude – wave’s height from the origin to the crest. • Wavelength (l)– distance between the crests. • Frequency (u)– number of wave cycles to pass a given point per unit of time.

  22. Water wave (ripple).

  23. Physics and the Quantum Mechanical Model • Frequency and wavelength are inversely proportional. As frequency increases, wavelength decreases, and vice versa, but their product will always equal the speed of light. • c = lu • SI units for frequency are cycles per second is a hertz (Hz), or 1/seconds (1/s or s-1).

  24. Relationship Between Wavelength and Frequency

  25. Physics and the Quantum Mechanical Model • What is the frequency of light that has a wavelength of 550 nm? (1m = 109 nm or 1 nm = 10-9 m)? • What is the wavelength of light, in cm, that has a frequency of 9.60 x 1014 Hz (1/s)? • What is the frequency of light (Hz) that has a wavelength of 740 nm (1m = 109 nm or 1 nm = 10-9 m)?

  26. Physics and the Quantum Mechanical Model • Sunlight splits into a spectrum of colors when it passes through a prism. • Colors of the spectrum include red, orange, yellow, green, blue, indigo and violet. • Red light has the longest wavelength and the lowest frequency, while violet light has the shortest wavelength and the highest frequency.

  27. Dispersion of White Light By a Prism

  28. A photon of red light (relatively long wavelength) carries less energy than a photon of blue light (relatively short wavelength) does.

  29. Physics and the Quantum Mechanical Model • Every element emits light after it absorbs energy. The light that is emitted (atomic emission spectra) is different for every element, and differs from white light because it is not continuous. • Max Planck said that color changes can be explained if you assume that the energy of a substance changes in small increments.

  30. Emission (line) Spectra of Some Elements

  31. Emission (line) Spectra of Some Elements (cont’d)

  32. Emmision (line) Spectra of Some Elements (cont’d)

  33. Physics and the Quantum Mechanical Model • Planck showed that the amount of radiant energy (E) absorbed or emitted by a substance is proportional to the frequency of the radiation. • E = hu • h is Planck’s constant (6.626 x 10-34 J s) • Any attempt to increase or decrease the energy of a system by a fraction of h times u will fail because energy is only emitted or absorbed in quanta, or bunches of energy.

  34. Planck’s Constant Examples • What is the energy of a photon with a frequency of 2.94 x 1015 cycles per second (s-1 or Hz)? • What is the energy of a light particle with a wavelength of 675 nm?

  35. Homework Problem Examples • What is the wavelength, in nm, of light with a frequency of 9.5 x 109 s-1? ( 1 m = 109 nm) • How much energy is contained in a photon with a wavelength of 5.17 x 10-4 m?

  36. 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 • The energy of the photon is directly proportional to the frequency of light • Higher frequency = More energy in photons

  37. Physics and the Quantum Mechanical Model • Photons – light energy. The energy of photons is quantized according to the equation E = hu. • Light was therefore thought to have a dual wave-particle behavior to explain all of its characteristics.

  38. Electromagnetic radiation (a beam of light) can be pictured in two ways: as a wave and as a stream of individual packets of energy called photons.

  39. Bohr’s Model • Energy of an electron 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

  40. Bohr’s Model • Atoms have a minimum energy called the ground state • 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

  41. (a) A sample of H atoms receives energy from an external source. (b) The excited atoms (H) can release the excess energy by emitting photons.

  42. When an excited H atom returns to a lower energy level, it emits a photon that contains the energy released by the atom.

  43. Hydrogen atoms have several excited-state energy levels.

  44. Each photon emitted by an excited hydrogen atom corresponds to a particular energy change in the hydrogen atom.

  45. 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. • 1st energy level can hold 2 electrons (e-1), the 2nd 8e-1, the 3rd 18e-1, etc. • farther from nucleus = more space = less repulsion

  46. Models of the Atom • Energy level – region around the nucleus where the electron is likely to be found. Think of steps on a ladder. • Essentially, you must be in one energy level or another, you can’t be between energy levels, just like you can’t stand in mid-air between the steps of a ladder.

  47. The difference between continuous and quantized energy levels can be illustrated by comparing a flight of stairs with a ramp.

  48. Models of the Atom • Energy levels are not equally spaced. The further away an electron is from the nucleus, the easier it becomes to pull that electron off of that particular atom. • Erwin Schrodinger – in 1926, he came up with a new way of describing the energy and location of an electron, called the quantum mechanical model, which is a mathematical method.

  49. Models of the Atom • The quantum mechanical model does not say that electrons take exact paths around the nucleus, but that it estimates the probability (likelihood) of finding an electron in a certain position. • If the electron cloud is very dense, it is more likely that you will find the electron there, then if the electron cloud is less dense.

  50. The probability map, or orbital, that describes the hydrogen electron in its lowest possible energy state.