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Chapter 5 The Periodic Law Section 5-1 History of the Periodic Table Stanislao Cannizzaro 13 Jul 1826 – 10 May 1910 Cannizzaro • Was the first scientist to accurately measure atomic masses. • This was incredibly important for the work that Dmitri Mendeleev was going to do. Dmitri Mendeleev 8 Feb 1834 – 2 Feb 1907 Dmitri Mendeleev • Arranged the elements on cards in order of increasing atomic mass • Found columns of elements with similar properties • There were gaps in his columns Basic Version of Mendeleev’s Periodic Table The gaps • Hypothesized the gaps were undiscovered elements. • Predicted the props of these elements. • Predicted them well. Mendeleev’s Mistake • There were irregularities when arranged according to atomic weight. Henry Mosely 23 Nov 1887 – 10 Aug 1915 Henry Mosely • Discovered a unique charge on the nucleus of the atom. • Arranged the elements according to increasing atomic number • When he did this, the irregularities disappeared. Periodic Law • The properties of elements tend to change with atomic number gradually, in a periodic way. John Strutt 12 Nov 1842 – 30 Jun 1919 William Ramsay 2 Oct 1852 – 23 Jul 1916 Strutt & Ramsay • In 1894, they discovered argon. • Nobody noticed it before because it is completely unreactive. • In 1868, helium had been discovered as part of the sun and in 1895, Ramsay showed its existence on earth. • In 1898, Ramsay discovered krypton and xenon. Friedrich Ernst Dorn discovered radon in 1900. Lanthanides • Discovered in the early 1900’s. • They are found in the f block • They are shiny and act like the alkaline earth metals. Actinides • All of these elements are radioactive. • They are found in the f block Alkali Metals • Group 1 • These elements are soft and can be cut with a knife. • They are highly reactive. The will react with both air and water. • They form alkaline/basic solutions (the opposite of acidic solutions). • Their electron configurations all end s1. Sodium Alkaline Earth Metals • Group 2 on the periodic table. • These elements are harder and denser than the alkali metals. • They are also reactive, but less so than the alkali metals. • They will also form alkaline/basic solutions. • Their electron configurations all end s2. Magnesium Hydrogen • This element doesn’t belong with any group. • Its electron configuration is 1s1. Helium • Even though its electron configuration ends s2, it isn’t an alkaline earth metal. It is a noble gas because its highest energy level orbitals are full. Transition metals • AKA transition elements • AKA d block elements • These elements are what we typically picture as common metals. • Their d orbitals are being filled. Transition metals cont’d • They are shiny and good conductors of electricity. • They are less reactive than the other metals. • Some like gold are highly unreactive. Chromium metal Main Group Elements • Properties of these elements vary greatly because they include metals, nonmetals, metalloids, and noble gases. • They include the elements of the s and p blocks. • Sometimes they are called the representative elements because metals, nonmetals, metalloids, and noble gases are all represented Halogens • These are the most reactive nonmetals. • They have 7 electrons in the outermost energy level and their electron configurations all end in s2p5. • They will react with metals to form salts. Noble Gases • • • • Group 18 No stable compounds for He, Ne, or Ar Very low reactivity for the rest Full s & p orbitals (s2p6) in the higest energy level – This is very stable - they have no need to react with anything else. Noble Gases cont’d • Most other atoms gain/lose e- to achieve this e- configuration • Ne & Ar are used in signs • He - Low density - Air ships & weather balloons Noble Gases Cont’d • Rn - Radioactive • Found in homes - Linked to Lung CA • Once you test for it, you must disclose the results to potential buyers 5 – 3 Electron Configuration and Periodic Properties Atomic Radius • As you move down a group it increases • The outermost e- are being added to higher energy levels (further from the nucleus. Atomic Radius • As you move across a period, it decreases • Even though e- are being added, they are added to the same energy level (same distance from the nucleus). Atomic Radius • The charge on the nucleus increases as you move across the period and so it has a “tighter” hold on the e- being added. Shielding Effect • The reduction of the attractive force between a nucleus and its outer electrons due to the blocking effect of inner electrons. Ionization Energy • The amount of energy needed to remove an electron from an atom Ion • An atom that has gained or lost an e• If it has gained an e-, it will be _____. • If it has lost an e-, it will be _____. Ionization Energy • As you go down a group, it decreases • Shielding effect and electrons are being added to higher energy levels. Ionization Energy • As you move across a period, it increases • The charge on the nucleus increases as you move across the period and so it has a “tighter” hold on the e- being added. Electron Affinity • Measures the tendency of an atom to attract electron • The energy change that occurs when an electron is acquired by a neutral atom. • Metals tend to have positive energy changes, they do not have a tendency to attract electrons • Non-metals tend to have negative energy changes, they have a strong tendency to attract electrons EA • As you move down a group, electrons add with greater difficulty (values become more positive) • Shielding Effect and electrons are being added to higher energy levels. • There are exceptions EA • Electrons add more easily as you move across the periodic table.(values become more neg) • The charge on the nucleus increases as you move across the period and so it has a “tighter” hold on the e- being added. Electronegativity • Tendency for an atom to attract e- to itself when combined with another atom. • F is the most EN • EN decreases as you move down a group • EN increases as you move across a period Electronegativity • Based on the Pauling Scale. • Linus Pauling • 28 Feb 1901 – 19 Aug 1994