Download
1 / 21
230 likes | 511 Vues
Chapter 5 Atomic Models. Different metal containing compounds emit colored light when fireworks burn. This colored light is a combination of a series of colors called a spectral pattern. ( 谱带 )
E N D
Chapter 5 Atomic Models Different metal containing compounds emit colored light when fireworks burn. This colored light is a combination of a series of colors called a spectral pattern. (谱带) Each element emits its own characteristic spectral pattern, which can be used to identify the element. This allowed the scientists in 1900s to develop models of the atom’s internal structure.
5.1 Models help us visualize the invisible world of atoms Atom is very small.
We can not see them in the usual sense. This is because light travels in waves and atoms are smaller than the wavelengths of visible light, which is the light that allows the human eye to see things. So we can not see atoms through the media of light, even with a microscope.
We can see atoms indirectly through scanning tunneling microscope (STM), which was invented in 1980s. Fig1:Scanning tunneling microscope Fig2:an image of gallium and arsenic atoms obtained with an STM fig3 : an STM image of mono-layer of perylene derivative on graphite substrate, where the epitaxial relationship is observed between the organic molecule and the substrate graphite
5.3 Atoms can be identified by the light they emit When we view the light from glowing atoms, we see that the light consists of a number of discrete frequencies rather than a continuous spectrum. This is called element’s atomic spectrum (原子光谱). In 1800s researchers noted the orderliness of element’s atomic spectrum, especially hydrogen, but could not give the explanation.
5.4 Niels Bohr used the quantum hypothesis to explain atomic spectra Max Planck’s quantum hypothesis (量子假设): a beam of light energy is not the continuous stream of energy, but consists of small, discrete packets of energy. Each packet was called a quantum. In 1905, Einstein recognized that these quanta of light behave like particles. Each quantum was called a photon (光子). Light behaves as both a wave and a particle.
Electron loses potential energy and moves closer to nucleus. A photon of light is emitted Bohr’s explanation Electron gains potential energy and moves farther from nucleus. A photon of light is absorbed
Bohr’s planetary model of atom There are only a limited number of permitted energy levels in an atom, and an electron can only stay in these energy levels. Each energy level has a principal quantum number n (主量子数). The energy level with n=1 has the lowest energy.
Photons are emitted by atoms as electrons move from higher-energy outer orbits to lower-energy inner orbits. The energy of an emitted photon is equal to the difference in energy between the two orbits. Because an electron is restricted to discrete orbits, only particular light frequencies are emitted.
5.4 Electrons exhibit wave properties Question by Louis de Groglie: If light has both wave properties and particle properties, why can not material particle, such as electron, also have both? Answer: Every particle of matter is somehow endowed with a wave to guide it as it travels. The more slowly an electron moves, the more its bahvior is that of a particle with mass. The more quickly it moves, the more its behavior is that of a wave of energy. In an atom, an electron moves at very high speeds: on the order of 2000000 meters per second. Practical application of the wave behavior of electrons: electron microscope.
The electron’s wave nature can be used to explain the Bohr’s planetary model: (1) Permitted energy levels are a natural consequence of electron waves closing in on themselves in a synchronized manner. (2) By viewing each electron orbit as a self-reinforcing wave, we that the circumference of the smallest orbit can be no smaller than a single wavelength.
Where 90% of the e- density is found for the 1s orbital e- density (1s orbital) falls off rapidly as distance from nucleus increases How to visualize electron waves? Probability clouds and atomic orbitals In 1926, Erwin Schrodinger formulated a equation from which the intensities of electron waves in an atom could be calculated. It was found that the intensity at any given location determined the probability of finding the electron at that location. The electron was mostly likely to be found where its wave intensity was grestest.
l = 0 (s orbitals) l = 1 (p orbitals) l = 2 (d orbitals) Atomic orbitals have shapes and sizes!
The size of orbital is indicated by Bohr’s principal quantum number n = 1, 2, 3, 4, 5, 6, 7… Fig5.19 The fluorine atom has five overlapping atomic orbitals that contain its nine electrons, which are not shown
5.6 Energy-level diagrams describe how orbitals are occupied Each orbital has a capacity of two, but no more than two, electrons. They spin in opposite directions, which generates two oppositely oriented magnetic fields that are attractive and partly compensate for the electrical repulsion between the electrons.
Example: Lithium (3): 1s22s1 Boron (5): 1s22s22p1 Carbon (6): 1s22s22p2 Nitrogen (7): 1s22s22p3 Oxygen (8): 1s22s22p4 The properties of an atom are determined mostly by its outermost electrons (最外层电子), since they are the ones in direct contact with the external environment.
The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule). C 6 electrons N 7 electrons O 8 electrons F 9 electrons Ne 10 electrons Ne 1s22s22p6 C 1s22s22p2 N 1s22s22p3 O 1s22s22p4 F 1s22s22p5 7.7
5.7 Orbitals of similar energies can be grouped into shells The seven rows correspond to the seven periods in the periodic table. The maximum number of electrons each row can hold is equal to the number of elements in the corresponding period. (2, 8, 8, 18, 18, 32, 32)
Increasing First Ionization Energy Increasing First Ionization Energy From electron configuration, we can predict: (1)The smallest atoms are at the upper right of the periodic table. The smallest atoms have the most strongly held electrons. (Represented by ionization energy (电离能), the amount of energy needed to pull an electron away from an atom). The ionization energy also determines the atom’s chemical behavior, which will be discussed in chapter 6.
