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Chapter 5 Structures of the Atom Electron Electrons were discovered in 1897 by Sir J. J. Thomas Electrons have the following characteristics: Travel from a cathode (electron emitter) in straight lines. Cause a piece of metal foil to become hot by striking it for a period of time. Deflected by magnetic and electric fields showing that they have a negative charge. Are emitted by some metals when heated. This is called “thermal emission”. Are emitted by some metals when exposed to light. This is called “photoelectric effect”. Proton Protons discovered in 1886 by Eugene Goldstein using a Crookes Tube Protons have the following characteristics: Positive charge same strength as an electron found in nucleus Neutron Neutrons were discovered in 1932 by James Chadwick Neutrons have the following characteristics: mass almost identical to proton neutral charge will disintegrate to form a proton and an electron when outside the nucleus Radioactivity and the Atomic Structure In 1896 Antoine Becquerel discovered that certain elements decomposed to form other elements by emitting radiation. There are three types of radiation alpha () particle which is a He+2 ion beta () particle which is an egamma () ray which is made up of high energy and has no charge Nuclear Atom In 1911 John Rutherford proposed that the atom was composed of a very small, positively charged nucleus where most of the mass is concentrated, surrounded by the number of electrons necessary to produce an electrically neutral atom. He was able to propose this by bombarding a thin gold foil with alpha particles. Because only a few alpha particles were deflected, Rutherford proposed that there must be a small nucleus within the atom. This is today’s model of the atom. Bohr Model of the Atom If electrons spin around the nucleus, why don’t they eventually slow down and get sucked into the nucleus? This was the question posed by many scientists. Niels Bohr proposed that the electrons are quantized, or restricted to discrete, individual areas or values. The energy of the electron is restricted to certain values which correspond to the radius of its orbit. The radius is characterized by an integer, n. The equation being: E kZ2 2 n where k = 2.179x10-18J; Z = the atomic number; and n = the integer characteristic of the orbit. Atomic Spectra and Atomic Structure An electron in an atom can be given enough energy to be promoted from one orbital to the next higher. When that electron moves back to its original orbital, that added energy is given off in the form of light. The amount of energy given off can be determined in the following equation where E is energy in Joules. 1 1 E 2.179 x1018 2 2 n1 n2 n1 is the lower-energy inner orbital n2 is the higher-energy outer orbital The wavelength of the emitted energy can be calculated by using the formula: 25 hc 1.98647x10 or E E The Quantum-Mechanical Model of the Atom Erwin Schrödinger showed that the electron may be visualized as being in rapid motion within one of several regions of space located around the nucleus. He called these regions “orbitals” Some scientists call the occupancy of these orbitals “electron densities” because of the rapidity of the electron moving within. The location of an electron cannot be determined exactly. We can identify only the region or volume of space where there is a relatively high probability of finding the electron Results of the Quantum-Mechanical Model of the Atom Orbitals are characterized by n, the principal quantum number, which may take on any integral value: n = 1, 2, etc. As n increases, the orbitals extend further away from the nucleus. A shell contains orbitals with the same value of n. A shell will have the values of K, L, M, etc. Shell K L M N n 1 2 3 4 Orbitals with the same value of n may have different shapes called the azimuthal quantum number distinguished by the second quantum number l. An electron with l = 0 will have a spherical orbital shape or s-orbital An electron with = 1 will be a p-orbital An electron with = 2 will be a d-orbital An electron with = 3 will be an f-orbital The shapes of the f-orbitals can be found in your text. So: Shell K L M N O n Value 1 2 3 4 5 Possible Values 0 0, 1 0, 1, 2 0, 1, 2, 3 0, 1, 2, 3, 4 Types of Orbibals 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g For values larger than zero, orbitals with the same value of in any given shell have the same shape and the same energy but differ in their orientation. There are 2 + 1 orbitals with the same quantum number that differ in their orientation Each orientation is characterized by a third quantum number, m, the magnetic quantum number. When = zero, there is only one orientation in space for the orbital When = 1, there are three possible values of m (-1, 0, +1) and thus the porbital has three sub-orbitals. When = 2, there are five possible values of m (-2, -1, 0, +1, +2) and thus the d-orbital has five sub-orbitals. Electrons have a fourth quantum number, s, called spin quantum number. The spin specifies the direction of spin of an electron about its own axis. It can be either clockwise or counterclockwise designated as s= +1/2 or s= -1/2 The energy of an electron in an atom is limited to discrete values. The energies can be determined from the wave function that describes the behavior of the electron. The maximum number of electrons that may be found in a shell with a principal quantum number of n is 2n2. The maximum number of atomic orbitals in a shell is given by n2. Orbital Energies and Atomic Structure The energies of atomic orbitals increase as the principal quantum number, n, increases. The energies of the orbitals increase within a shell as the quantum number, , increases. For atomic numbers greater than 20, the relative energies of the orbitals may differ slightly from the order shown. For example, the energies of the electrons in 1s orbitals become lower and lower as the atomic numbers of the atoms increase. It is possible to predict the electron configuration of an atom in its ground state using the following guidelines: Electrons occupy the lowest possible energy levels or orbitals The Pauli Exclusion Principle states that no two electrons in the same atom can have the same set of four quantum numbers. Hund’s Rule: Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin. The Aufbau Process The systematic variations in the electronic structures of the various elements can be “built” in atomic order. This process is called the Aufbau process Note the exceptions to the aufbau process found in your text book. Some of the exceptions are chromium, copper and lanthanum. These exceptions are usually due to very similar energies within the subshells.