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Why are Electrons Important? C hemistry Unit 4

Why are Electrons Important? C hemistry Unit 4

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Why are Electrons Important? C hemistry Unit 4

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  1. Why are Electrons Important?Chemistry Unit 4

  2. Why are Electrons Important?

  3. Main Ideas • Light, a form of electronic radiation, has characteristics of both a wave and a particle • Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals. • Atomic emission spectra corresponds to the release of energy from an electron changing atomic energy levels. • A set of three rules determines the arrangement in an atom.

  4. 4:1 Light and Quantized Energy

  5. Light and Quantized Energy Objectives: Compare the wave and particle natures of light. 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.

  6. 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.

  7. The Atom and Unanswered Questions 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. In order to understand this relationship and the nature of atomic structure, it will be helpful to first understand the nature of light.

  8. Wave Nature of Light Electromagnetic radiationis a form of energy that exhibits wave-like behavior as it travels through space. Visible light Microwaves X-rays Radio waves

  9. Wave Nature of Light All waves can be described by several characteristics. The wavelength(λ) is the shortest distance between equivalent points on a continuous wave. (crest to crest, trough to trough) • The frequency(ν) is the number of waves that pass a given point per second. • Hertz- SI unit for frequency= one wave/sec • Energy increases with increasing frequency

  10. Wave Nature of Light All waves can be described by several characteristics. • The amplitude is the wave’s height from the origin to a crest. • Independent of wavelength and frequency

  11. Wave Nature of Light

  12. Wave Nature of Light The speed of light (3.00 x108 m/s) is the product of it’s wavelength and frequency c = λν.

  13. Wave Nature of Light All electromagnetic waves, including visible light travels at 3.00 x 108m/s in a vacuum. Speed is constant but wavelengths and frequencies vary. • Sunlight contains a continuous range of wavelengths and frequencies. • A prism separates sunlight into a continuous spectrum of colors.

  14. Wave Nature of Light • The electromagnetic spectrum includes all forms of electromagnetic radiation. • Not just visible light.

  15. Wave Nature of Light

  16. 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. A quantum is the minimum amount of energy that can be gained or lost by an atom.

  17. Particle Nature of Light • Max Planck (1858-1947) – matter can gain or lose energy only in small amounts. • E=hv • Planck’s constanthas a value of 6.626 x10–34 J ● s. • Energy can only be emitted or absorbed in whole number multiples of h.

  18. Particle Nature of Light The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.

  19. Particle Nature of Light Albert Einstein proposed in 1905 that light has a dual nature. Nobel prize in 1921. A beam of light has wavelike and particlelike properties. A photon is a mass-less particle of electromagnetic radiation with no mass that carries a quantum of energy. Ephoton = hνEphoton represents energy.h is Planck's constant.νrepresents frequency.

  20. Example Problem 1 Microwaves are used to cook food and transmit information. What is the wavelength of a microwave that has a frequency of 3.44 x 109 Hz?

  21. Example Problem 2 Microwaves are used to cook food and transmit information. What is the wavelength of a microwave that has a frequency of 3.44 x 109 Hz? What is the amount of energy in this wavelength?

  22. Example Problem 2 While an FM radio station broadcasts at a frequency of 94.7 MHz, an AM station broadcasts at a frequency of 820 kHz. What are the wavelengths of the two broadcasts?

  23. Example Problem 2 While an FM radio station broadcasts at a frequency of 94.7 MHz, an AM station broadcasts at a frequency of 820 kHz. What are the wavelengths of the two broadcasts? What is the associated energy for each of these broadcasts?

  24. Example Problem 3 Every object gets its color by reflecting a certain portion of visible light. The color is determined by the wavelength of the reflected photons, and therefore their energy. The blue color in some fireworks occurs when copper (I) chloride is heated to approximately 1500K and emits blue light of wavelength 4.50 x 102 nm. How much energy does one photon of this light carry?

  25. Atomic Emission Spectrum The atomic emission spectrumof an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element. • Emission lines are specific to an element and can be used for identification.

  26. Atomic Emission Spectrum

  27. Absorption Spectra The absorption spectra of an element is the set of frequencies of the electromagnetic waves absorbed by the atoms of the element. • Absorption lines are specific to an element and can be used for identification.

  28. Emission vs. Absorption

  29. Question? What is the smallest amount of energy that can be gained or lost by an atom? A. electromagnetic photon B. beta particle C. quanta D. wave-particle

  30. Question? What is a particle of electromagnetic radiation with no mass called? A. beta particle B. alpha particle C. quanta D. photon

  31. 4:2 Quantum Theory of the Atom

  32. Quantum Theory of the Atom Objectives: Compare the Bohr and quantum mechanical models of the atom. Explain the impact of de Broglie's wave article duality and the Heisenberg uncertainty principle on the current view of electrons in atoms. Identifythe relationships among a hydrogen atom's energy levels, sublevels, and atomic orbitals.

  33. Bohr’s Model of the Atom Bohr correctly predicted the frequency lines in hydrogen’s atomic emission spectrum. The lowest allowable energy state of an atom is called its ground state. When an atom gains energy, it is in an excited state.

  34. Bohr’s Model of the Atom Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits. The smaller the electron’s orbit the lower the atom’s energy state or level

  35. Bohr’s Model of the Atom Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits. The larger the electron’s orbit the higher the atom’s energy state or level.

  36. Bohr’s Model of the Atom Each orbit was given a number, called the quantum number. The orbit closest to the nucleus is n=1

  37. Bohr’s Model of the Atom Example: Hydrogen’s single electron is in the n = 1 orbit in the ground state. Atom does not radiate energy. When energy is added, the electron moves to the n = 2 orbit. Atom is excited. When electron moves from an excited state to ground state, a photon is emitted.

  38. Electron States

  39. Quantum Mechanical Model The Quantum Mechanical Model of the Atom – this model progressed through a series of scientific findings: • Louis de Broglie (1892–1987) hypothesized that particles, including electrons, could also have wavelike behaviors. • Like vibrating guitar strings – multiples of half waves. • Orbiting electron – whole number of wavelengths.

  40. Quantum Mechanical Model

  41. λrepresents wavelengthsh is Planck's constant.m represents mass of the particle.vrepresents velocity. Quantum Mechanical Model The de Broglie equationpredicts that all moving particles have wave characteristics.

  42. Example Problem Why do we not notice the wavelengths of moving objects such as automobiles?

  43. Quantum Mechanical Model Heisenberg showed it is impossible to take any measurement of an object without disturbing it. • The Heisenberg uncertainty principlestates that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time. • Means that it is impossible to assign fixed paths for electrons like the circular orbits as previously thought.

  44. Quantum Mechanical Model Heisenberg showed it is impossible to take any measurement of an object without disturbing it. • The Heisenberg uncertainty principlestates that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time. • The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.

  45. Quantum Mechanical Model Schrödinger expanded on the de Broglie wave particle theory and created the quantum mechanical model that we know today. • treated electrons as waves in a model called the quantum mechanical model of the atom. Schrödinger’s equation applied equally well to elements other than hydrogen.

  46. Quantum Mechanical Model Both models limit an electron’s energy to certain values. Unlike the Bohr model, the quantum mechanical model makes no attempt to describe the electron’s path around the nucleus. Electrons are located around the nucleus at a position that can be described only by a probability map. A boundary surface is chosen to contain the region that the electron can be expected to occupy 90% of the time.

  47. Quantum Mechanical Model The wave function predicts a three-dimensional region around the nucleus called the atomic orbital.

  48. Quantum Numbers and the Revised Model The revised model defines the relationship between an electron’s energy level, sublevel and atomic orbitals. • Four quantum numbers make up the identification of each electron in an atom.

  49. Atomic Orbitals Principal quantum number (n) indicates the relative size and energy of atomic orbitals. n specifies the atom’s major energy levels, called the principal energy levels.

  50. Atomic Orbitals Energy sublevels (s,p,d or f) are contained within the principal energy levels.