Chapter 10. Modern Atomic Theory and the Periodic Table

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1 Chapter 10 Modern Atomic Theory and the Periodic Table 1

2 10.1 A brief history 10.1 A brief history atoms proposed by Greek philosopher Dalton s model of atom Thomson s model Rutherford s model there remain questions can not be answered: how atomic structure relates to the periodic table (arrangement of electrons) how to explain line spectrum of atom 2

3 10.2 Electromagnetic radiation energy and light energy travels through space is by electromagnetic radiation, all radiations travel at the same speed v = λ υ = m/s λ : wavelength υ : frequency electromagnetic spectrum wavelike nature -- radiation behaves like particle -- photon explain the properties of electromagnetic radiation by both wave and particle properties 3

4 10.3 The Bohr atom at high temperature or when subjected to high voltage, elements in the gaseous state give off colored light a set of brightly colored lines line spectrum ex. line spectrum of hydrogen line spectrum indicates that light is being emitted only at certain wavelength (or frequency) 1912 Bohr model of hydrogen electrons exit in specific regions at various distance from the nucleus the electrons as revolving in orbits around the nucleus like planets rotating around the Sun 4

5 Max Planck energy quanta the energy is never emitted in a continuous stream but only in small discrete packets called quanta electrons are only in several energy levels hydrogen atom absorbs one or more quanta of energy, the electron will jump to a higher energy level ground state the lowest energy level excited states the higher energy levels when an electron falls from a high-energy level to a lower one, a quantum of energy is emitted as light at a specific frequency Bohr model 1) suggesting quantized energy levels for electrons 2) showing that spectral lines result from the radiation of small increments of energy when electrons shift from one energy level to another however, Bohr model only succeeded in H atom, 5 did not succeed for heavier atoms

6 1924 de Broglie all objects have wave properties for small objects such as an electron, the wave properties become significant 1926 Schrödinger Schrödinger equation a mathematical model that described electrons as waves the probability of finding an electron in a certain region around the atom can be determined wave mechanics or quantum mechanics forming the basis for our modern understanding of atomic structure we cannot locate an electron precisely within an atom electrons are not revolving around nucleus in orbits as Bohr postulated orbital a region where a high probability of finding a given electron 6

7 10.4 Energy levels of electron Bohr the energy of the electron is quantized the electron is restricted to only certain allowed energies wave-mechanics also predicts discrete principal energy levels within the atom these energy levels are designated by the letter n n: positive integer as n increases, the energy of the electron increases, and the electron is found on average farther from the nucleus each principal energy level is divided into sublevels n = 1 n = 2 n = 3 1 sublevel 2 sublevels 3 sublevels each of these sublevels contains space for electrons called orbitals 7

8 n = 1 1s orbital spherical shape the electron does not move around on the surface of sphere, the surface encloses a space where there is a 90% probability that the electron may be found how many electrons can fit into a 1s orbital? spin a property of electron each electron can only spin in two directions representation of the spin or two electrons with the same spin cannot occupy the same orbital Pauli exclusion principle an atomic orbital can hold a maximum of two electrons which must have opposite spins n = 1 energy level contains one type of orbital (1s) that hold a maximum of 2 electrons n = 2 2s spherical shape hold a maximum of two electrons 2p 2p x, 2p y, 2p z 8

9 each p orbital has two lobes and can hold a maximum of two electrons the total number of electrons that can reside in all three p orbitals is 6 n = 3 n = 2 energy level contains two types of orbitals (a 2s and three 2p) that hold a maximum of 8 electrons 3s 3p 3p x, 3p y, 3p z 3d 3d xz, 3d xy, 3d yz, 3d z 2, 3d x 2 y 2 n = 3 energy level contains three types of orbitals (a 3s, three 3p, five 3d) that hold a maximum of 18 electrons 9

10 n = 4 energy level contains four types of orbitals (a 4s, three 4p, five 4d, seven 4f) that hold a maximum of 328 electrons hydrogen atom consists of a nucleus (one proton) and one electron occupying a region outside of the nucleus cm 10-8 cm 10

11 10.5 Atomic structures of the first 18 elements all atoms contain orbitals similar to those found in hydrogen systematically placing electrons in these hydrogen like orbitals, following the guidelines: 1. no more than two electrons can occupy one orbital 2. electrons occupy the lowest energy orbitals available s < p < d < f for a given n value 3. each orbital on a sublevel is occupied by a single electron before a second electron enters ex. atomic structure diagrams of F, Na and Mg there are two ways to show the arrangement of the electrons in the orbitals: i) electron configuration number of electrons in sublevel orbitals 2p 6 principal energy level ii) orbital diagram orbital type of orbital spin 11

12 valence electrons the electrons un the outmost (highest) energy level of an atom ex. O 1s 2 2s 2 2p 4 6 valence electrons 12 Mg 1s 2 2s 2 2p 6 3s 2 2 valence electrons

13 10.6 Electron structures and the periodic table 1869 Mendeleev & Meyer periodic arrangements of the elements based on increasing atomic masses periodic table period horizontal row the outmost energy level group or family vertical column elements behave in a similar manner IA ~ VIIA, IB ~ VIIB, VIII, noble gases 1 ~ 18 representative elements A group elements transition elements B group elements IA alkali metals IIA alkaline earth metals 13 VIIA halogens

14 the valence electron configurations for H ~ Ar the valence electron configuration for the element in each column is the same, but the number for the energy level is different the chemical behavior and properties of elements in a particular family are similar and must be associated with the electron configuration abbreviated electron configuration B 1s 2 2s 2 2p 1 [He] 2s 2 2p 1 Cl 1s 2 2s 2 2p 6 3s 2 3p 5 [Ne] 3s 2 3p 5 Na 1s 2 2s 2 2p 6 3s 1 [Ne] 3s 1 n = 4 K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 [Ar] 4s 1 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar] 4s 2 element 21 ~ 30 transition elements electrons are placed in the 3d orbitals 14

15 arrangement of elements according to the sublevel being filled inner transition elements lanthanide series 4f actinide series 5f ex write the electron configuration for P and Sn P 1s 2 2s 2 2p 6 3s 2 3p 3 [Ne]3s 2 3p 3 Sn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 2 [Kr] 5s 2 4d 10 5p 2 15

16 groups of elements show similar chemical properties because of the similarity of these outmost electron configurations the periodic table illustrates several important points: 1. the number of the period corresponds with the highest energy level occupied by electrons 2. the group numbers for the representative elements are equal to the total number of outmost electrons in the atom 3. the elements of a family have the same outmost electron configuration, but in different energy level 4. the elements within each of the s, p, d, f blocks are filling the s, p, d, f orbitals 5. within the transition elements some 16 discrepancies in the order of filling occur

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