Chapter 5 - The Periodic Law 5-1 History of the Periodic Table I. Mendeleev's Periodic Table (1869) A. Organization by rows and columns (the reverse of what we have today) 1. placed elements in columns in order of atomic weight a. made some exceptions to place elements in rows with similar properties *Tellurium and iodine's places were switched 2. horizontal rows had elements with similar chemical properties B. Missing Elements: 1. gaps existed and he predicted the properties of the ?yet to be discovered? elements * Scandium, germanium and gallium: 3 elements he predicted C. Unanswered Questions: 1. Why didn't some elements fit in order of increasing atomic mass? 2. Why did elements exhibit periodic behavior? II. Moseley and the Periodic Table (1911) A. Protons and Atomic Number: x-ray experiments used to determine # of protons * the periodic table was found to be in atomic number order, not atomic mass order *the tellurium-iodine anomaly was explained B. The Periodic Law: physical and chemical properties of elements are periodic functions of their atomic #s III. The Modern Periodic Table A. Discovery of Noble Gases: (most important since all atoms want to be like Noble Gases) 1. Noble Gases discovered late due to their chemical unreactivity 2. 1868 - Helium discovered as a component of the sun, based on the emission spectrum of sunlight 3. Ramsey discovered helium (on earth), argon, krypton, and xenon and placed new group between Halogens and Alkali Metals B. Periodicity: elements with similar properties are found at regular intervals within the "periodic" table 5-2 Electron Configuration and the Periodic Table I. Periods and the Blocks of the Periodic Table A. Periods: horizontal rows on the periodic table * period number corresponds to the highest principal quantum number of the elements in the period B. Sublevel Blocks: table can be broken into blocks corresponding to s, p, d, f sublevels II. Blocks and Groups A. s-Block: Groups 1 and 2 1. Group 1: The Alkali Metals: one s electron in outer shell a. Soft, silvery metals of low density and low melting points b. Highly reactive, never found pure in nature 2. Group 2: The Alkaline Earth Metals: two s electrons in outer shell a. Denser, harder, stronger, less reactive than Group 1 b. Too reactive to be found pure in nature B. d-Block: Groups 3 ? 12: Transition Metals 1. metals with typical metallic properties (luster, malleability, ductility, conductivity) 2. Group number = sum of outermost s and d electrons C. p-Block elements: Groups 13 - 18 1. Properties vary greatly a. Metals: (1) softer and less dense than d-block metals (2) harder and more dense than s-block metals b. Metalloids: (1) Brittle solids with some metallic and nonmetallic properties (2) Semiconductors c. Nonmetals: (1) Halogens (Group 17) are most reactive of the nonmetals D. f-Block: Lanthanides and Actinides 1. Lanthanides are shiny metals similar in reactivity to the Group 2 metals 2. Actinides: a. All are radioactive b. Plutonium (94) through Lawrencium (103) are man-made 5-3 Electron Configuration and Periodic Properties I. Atomic Radii: A. Atomic Radius: One half the distance between nuclei of identical atoms bonded together B. Trends: 1. Atomic radius tends to decrease across a period due to increasing positive nuclear charge 2. Atomic radii tend to increase down a group due to increasing number energy levels (outer electrons are farther from the nucleus) II. Ionization Energy: A. Ion: An atom or a group of atoms that has a positive or negative charge B. Ionization: Any process that results in the formation of an ion C. Ionization Energy: The energy required to remove one electron from a neutral atom of an element forming a cation, measured in kilojoules/mole (kJ/mol) A + energy --> A+ + e- (?A? represents symbol of any element) D. Trends: 1. Ionization energy of main-group elements tends to increase across each period a. Atoms are getting smaller, electrons are closer to the nucleus 2. Ionization energy of main-group elements tends to decrease down a group a. Atoms are getting larger, electrons are farther from the nucleus b. Outer electrons become increasingly more shielded from nucleus by inner e- 3. Metals have a characteristic low ionization energy 4. Nonmetals have a high ionization energy 5. Noble gases have a very high ionization energy E. Removing Additional Electrons Na + 496 kJ/mol --> Na+ + e- Na+ + 4562 kJ/mol --> Na++ + e- Na++ + 6912 kJ/mol --> Na+++ + e- 1. Ionization energy increases for each successive electron 2. Each electron removed experiences a stronger effective nuclear charge 3. The greatest increase in ionization energy comes when trying to remove an electron from a stable, noble gas configuration III. Electron Affinity A. Electron Affinity: The energy change that occurs when an electron is acquired by a neutral, atom, measured in kJ/mol a. Most atoms release energy when they acquire an electron A + e- --> A- + energy (exothermic) b. Some atoms must be forced to gain an electron A + e- + energy --> A- (endothermic) B. Trends: 1. Halogens have the highest electron affinities 2. Metals have characteristically low electron affinities 3. Affinity tends to increase across a period 4. Electron affinity tends to decrease down a group 5. Second electron affinities are always positive (endothermic) IV. Ionic Radii A. Cations: Positive ions which are always smaller than corresponding atom 1. Protons outnumber electrons 2. Less shielding of electrons B. Anions: Negative ions which are always larger than corresponding atom 1. Electrons outnumber protons 2. Greater electron-electron repulsion C. Trends: 1. Ion size tends to increase downward within a group (same as atomic radii) V. Valence Electrons A. Valence Electrons: The electrons available to be lost, gained, or shared in the formation of chemical compounds * Main group element valence electrons are outermost energy level s and p sublevels Group #: 1 2 13 14 15 16 17 18 # of Valence Electrons: 1 2 3 4 5 6 7 8 VI. Electronegativity A. Electronegativity: A measure of the ability of an atom in a chemical compound to attract electrons * Elements that do not form compounds are not assigned electronegativities B. Trends: 1. Nonmetals have characteristically high electronegativity a. Highest in the upper right corner 2. Metals have characteristically low electronegativity a. Lowest in the lower left corner of the table 3. Electronegativity tends to increase across a period 4. Electronegativity tends to decrease down a group of main-group elements
Want to see the other 4 page(s) in Chapter_5_Notes.doc?JOIN TODAY FOR FREE!