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The Periodic Table Why is the Periodic Table important to me? • The periodic table is the most useful tool to a chemist. • You get to use it on every test. • It organizes lots of information about all the known elements. Pre-Periodic Table Chemistry … • …was a mess!!! • No organization of elements. • Imagine going to a grocery store with no organization!! • Difficult to find information. • Chemistry didn’t make sense. Dmitri Mendeleev: Father of the Periodic Table SOME PROBLEMS… HOW HIS WORKED… • Put elements in rows by • He left blank spaces for what he said were increasing atomic weight. undiscovered elements. • Put elements in columns (Turned out he was by the way they reacted. right!) • He broke the pattern of increasing atomic weight to keep similar reacting elements together. Periodic Table • This table is a remarkable way to show the manifold relationships between differing kinds of elements • The modern table was devised in 1869 by Dimitri Mendeleyev • He arranged the elements by weight and by their chemical properties Periodic Table • "...if all the elements be arranged in order of their atomic weights a periodic repetition of properties is obtained." - Mendeleyev The Current Periodic Table • Mendeleev wasn’t too far off. • Now the elements are put in rows by increasing • ATOMIC NUMBER!! This is called “periodic law”. • The horizontal rows are called periods and are labeled from 1 to 7. • The vertical columns are called groups are labeled from 1 to 18. Metals vs Non Metals Metals • • • • • • • Have Luster (shiny / reflect light) Solid at room temperature (except mercury) Gray or silver color (except gold and copper) Malleable (Bend, dent, deform when hit with mallet) Ductile (Can be drawn into thin wire) Good conductor of electricity Good conductor of heat METALS Metals vs Non Metals NON Metals • Are dull (do NOT reflect light) • Some are solid, liquid or gas at room temperature • Come in many colors including colorless, black, purple, yellow, green, gray, etc. • Brittle (shatter when struck with mallet) • Poor conductor of electricity • Poor conductor of heat NONMETALS Metalloids or semi-metals Border the staircase Have blends of properties: some metal characteristics some non metal characteristics Si, Ge, As, Sb, Te, sometimes B METALLOIDS B Si Ge As Sb Te Periods and Groups • Horizontal rows in the periodic table are called periods. • Vertical columns are called groups. • We know that the electron configuration of the elements matches the way they are arranged. (see p362-363 in your textbook) The relation between orbital filling and the periodic table Groups…Here’s Where the Periodic Table Gets Useful!! • Elements in the same group have similar chemical and increasing physical properties!! • (Mendeleev did that on purpose.) Why?? • They have the same number of valence electrons. • They will form the same kinds of ions. • They increase in size from smallest at top to largest at bottom. Families on the Periodic Table • Columns are also grouped into families. • Families may be one column, or several columns put together. • Families have names rather than numbers. • Most noted families of the periodic table are A separate slide for each follows • Alkali Metals • Alkaline Earth Metals • Halogens • Noble Gases Hydrogen • Hydrogen is an exception, as it 1 s sometimes is shown with the alkali metals family because of its electron configuration and sometimes with the halogens because it needs one more to fill its valence shell; it belongs to a family of its own. • Hydrogen is a diatomic element and a reactive gas. • In water, hydrogen will usually make the +1 ion, but can make the -1 ion. Alkali Metals s1 • 1st column on the periodic table (Group 1) not including hydrogen. • Very reactive metals, always combined with something else in nature (like in salt). • Soft enough to cut with a butter knife. • Makes the +1 ion. Alkaline Earth Metals s2 • Second column on the periodic table. (Group 2) • Reactive metals that are always combined with nonmetals in nature. • Several of these elements are important mineral nutrients (such as Mg and Ca. • Make the +2 ion. Transition Metals s2 d1through10 • Elements in groups 3-12, the d-block. • Less reactive harder metals • Includes metals used in jewelry and construction. • Metals used “as metal.” • f-block is called the inner transition metals. s2 f1through14 • Sometimes can make more than one ion: called polyvalent. Boron Family s2 p1 • Elements in group 13 • Aluminum metal was once rare and expensive, not a “disposable metal.” • This family crosses the metal /nonmetal line, its family pattern is not always apparent. Carbon Family s2 p2 • Elements in group 14 • Contains elements important to life and computers. • Carbon is the basis for an entire branch of chemistry. • Silicon and Germanium are important semiconductors. Nitrogen Family s2 p3 • Elements in group 15 • Nitrogen makes up over ¾ of the atmosphere. • Nitrogen and phosphorus are both important in living things. • The red stuff on the tip of matches is phosphorus. • Typically forms the -3 ion in ionic compounds. Oxygen Family or Chalcogens s2 p4 • Elements in group 16 • Oxygen is necessary for respiration. • Many things that stink, contain sulfur (rotten eggs, garlic, skunks,etc.) • When in an ionic compound, the -2 ion usually forms. Halogens s2 p5 • Elements in group 17 • Very reactive, volatile, diatomic, nonmetals • Always found combined with other element in nature . • Used as disinfectants and to strengthen teeth. • When in an ionic compound, halogens form the -1 ion. The Noble Gases s2 p6 • Elements in group 18 • VERY unreactive, inert, monatomic gases • Used in lighted “neon” signs • Used in blimps to fix the Hindenberg problem. • Have a full valence shell. • Does not bond. • Does not form an ion. • Group 1 is called the alkali metals and does not include hydrogen. • Group 2 is called the alkaline earth metals. • Groups 3-12 are called the transition metals • Groups 13-15 are not usually named • Group 16 are the chalcogens or oxygen family. • Group 17 are the halogens • Group 18 are the noble gases & are basically inert To review: • • • • • • • • Valence Shell by Group #: Group 1 Group 2 Groups 13 Groups 14 Groups 15 Group 16 Group 17 Group 18 s1 s2 s2p1 s2p2 s2p3 s2p4 s2p5 s2p6 These groups are called the “representative elements”. The representative elements in the s-block and p-block are more predictable and have less exceptions than the d-block and f-block elements. You must know these to be successful in Unit 4 (e- configuration and periodic table) and Unit 5 (bonding) • • • • • • • • Group 1 Group 2 Groups 13 Groups 14 Groups 15 Group 16 Group 17 Group 18 s1 s2 s2p1 s2p2 s2p3 s2p4 s2p5 s2p6 These are the valence shell electron configurations that we will be using repeatedly. Write them down. Study them. Know how they are used. Periods and Groups • Sizes of the atoms increase as we move from top to bottom down a group. • This is due to the increasing distance the electrons are from the nucleus. Each time a new principle energy level (shell) is added the orbiting electrons are at a higher energy and are further away from the nucleus. Periods and Groups • Sizes of the atoms decrease as we move from left to right across a period. • This is due to the increasing number of protons in the nucleus, so the electrical attraction between the nucleus and the orbiting electrons gets stronger and pulls the electrons closer to the nucleus. • Trends within periods and groups – There is usually a trend within a period (across, horizontal) and a trend within a group (down, vertical) • Atomic Radius trend – DECREASES ACROSS the period, INCREASES DOWN the group… Atomic Radius • Atomic size is a periodic trend influenced by electron configuration • Electron clouds do NOT have a clearly definded edge. • The outer limit of an electron cloud is defined as a spherical surface within which there is a 90% probability of finding an electron. • Atomic size is defined by how closely an atom lies to a neighboring atom. Electronegativity • We can display the table to demonstrate other properties as well • As you move from left to right across a period, the ability of the atom to attract electrons in a shared bond increases • This property is called electronegativity Electronegativity • The Electronegativity of an element is the relative ability of it’s atoms to attract electrons in a chemical bond. • Trends – INCREASES as you go ACROSS THE PERIOD, INCREASES as you go UP THE GROUP. Ionization Energy • Ionization energy is defined as the energy required to remove an electron from a gaseous atom…making the element a positive ion • Ionization energy can give us an indication of how strongly an atom’s nucleus holds onto its valence electrons. More about valence electrons later • High Ionization energy values = nuclei’s with strong holds on its electrons, which, in return, Ionization Energy’s Periodic Trend • Within the PERIOD, ionization energy INCREASES as you ACROSS the row. • Within the GROUP, ionization energy INCREASES as you go UP the column. Periodicity • 3 Kinds of Periodic Trends… • General trends are called periodicity. The properties repeat in a periodic pattern. • Atomic Radius, Electronegativity, Ionization Energy, Metallic Character, Ionic Radius and other properties, are periodic trends. Know the reasons as well as the trends: properties related to electron configuration TRENDS, TRENDS, TRENDS! Periodicity • Left to right horizontal/across a period: – greater positive charge of nucleus: • Holds/attracts electrons in the SAME energy level more tightly. • Top to bottom vertical in a group: – more energy levels, farther distance from nucleus: • Attractive force is reduced, electrons at HIGHER energy levels more loosely held. EXCEPTIONS: bumps and blips in • trends Metals to nonmetal trend often dictates type of ion formed: positive cation vs. negative anion. • Full sublevels are slightly more stable than partially filled sublevels: s2 more stable than s1 • Half filled sublevels are slightly more stable than partially filled sublevels: p3 more stable than p4 Both full s and p sublevels (s2p6 most stable, like noble gases) have a “sheilding effect” dropping attractive force on valence electrons. Abbreviated noble gas electron configuration helps in recognizing/remembering this effect. To Review • Group 1 is called the alkali metals and does not include hydrogen. • Group 2 is called the alkaline earth metals. • Groups 3-12 are called the transition metals • Groups 13-15 are not usually named • Group 16 are the chalcogens • Group 17 are the halogens • Group 18 are the noble gases & are basically inert Lots of Questions • Why are there only two elements in the first period? • Why are there eight elements in the second and third periods? • Why do we jump to eighteen elements across in the fourth period? • Why are the chemical properties of a group so similar? Lots of Questions • What you learned about electron configuration should answer the first three questions. • Our next unit on bonding will address the last question. • This chapter will involve discovering the periodic trends, making predictions based on the trends and deducing the reasons behind this periodictity • Read Chapter 14 for next week!