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The Periodic Table Concepts to Master • • • • • • • • • • • • • • • Who was important in the development of the periodic table and why? What is the difference between a chemical and physical property? What are elements arranged on the periodic table horizontally? What are elements arranged on the periodic table vertically? How many electrons can each orbital contain? Why can the periodic table be subdivided into s,p,d, and f blocks? What are the other names for the s, p, d, and f blocks? Is atomic radius a chemical or physical property? Is electronegativity a chemical or physical property? Why are the noble gases inert? Why do the alkali metals react so readily with the halogens? Why do the transition metals have multiple oxidation numbers? What is the trend in metallic character as you go down the periodic table? What is the common theme as you go down the periodic table? What is the diagonal rule? • • • • • • • • • • • • • • Where are electrons located? What are the characteristics of metals? What are the characteristics of nonmetals? What are the characteristics of metalloids? Explain periodicity? How are sublevels and PELs related? How are orbitals and orbits related? What’s the formula for the maximum number of electrons allowed in a certain PEL? List the elements that exist as diatomics. What are the trends in atomic radius as you go down and across the periodic table? Why? What are the trends in electronegativity as you go down and across the periodic table? Why? What are the trends in ionization energy as you go down and across the periodic table? Why? What is the most reactive metal and the most reactive non-metal? What is the common theme as you go across the periodic table? Vocab • • • • • • • • • • • • • • • • • • • Alkali Metals Alkaline Earth Metals Anion Atom Atomic Radius Boiling point Cation Density Diatomic Diatomic Ductile Electronegativity Groups Halogens Inert Gases Ionization Energy Kernel Lustrous Malleable • • • • • • • • • • • • • • • • • • • Melting point Metalloids Metals non metals Modern Periodic Law Monatomic Orbitals Orbits Periods Periodic Law Polyatomic Principle energy level Principle quantum number Reactivity Reactivity Sublevels Transition Metals Valence electrons Valence shell Labs: Graphing Trends and Constructing a Table Cool Websites • http://www.colorado.edu/physics/2000/app lets/a2.html • http://www.youtube.com/watch?v=DYW50 F42ss8 – Element song People and Periodic Table • Dmitri Mendeleev – Credited with organization of FIRST periodic table – Mendeleev’s greatest achievement was recognizing the fundamental rule that the chemical elements show an approximate repetition in their properties. • Elements were arranged by increasing atomic mass. • Elements were listed in columns so that those with similar properties were side by side. • He predicted the existence and properties of new elements (blank spaces in the first periodic table). People and Periodic Table • Henry Moseley – With the discovery of isotopes of the elements, it became apparent that atomic mass was not the significant player in the periodic law as Mendeleev, had proposed. – Moseley used X-rays to determine the atomic number of the known elements and then arranged them according to increasing atomic number. – Because of Moseley's work, the modern periodic table is based on the atomic numbers of the elements. pg186 Mosely • • • • • Student in Rutherford ‘s lab at the univ of Manchester Remember Rutherford was playing with radioactive sources to do gold-foil exp so he had access to xrays Found a mathematical relationship between the amt of energy in the xray beam (wavelength) and the # of protons in nucleus. This was huge because at the time Rutherford’s idea of a nucleus wasn’t “proven” (not enough scientists had repeated it) so it was too tentative for other scientists to accept. Mosely’s work confirmed Rutherfords nucleus conclusion . – – – • His work could be repeated by anyone Linked the order of the elements with a physical characteristic based on the atoms structure Now scientists knew what to look for when searching for new elements He died in battle during WWI – he was 27 Periodic Law • Mendeleev - The properties of the elements are a periodic function of their atomic masses. • Moseley - The properties of the elements are a periodic function of their atomic numbers. • Modern Periodic Law states that many of the physical and chemical properties of the elements tend to recur in a systematic manner with increasing atomic number. – Periods are the horizontal rows in the table. – Progressing from the lightest to the heaviest atoms, certain properties of the elements approximate those of precursors at regular intervals of 2, 8, 18, and 32 (periodicity). – Examples: • The 2d element (helium) is similar in its chemical behavior to the 10th (neon), as well as to the 18th (argon), the 36th (krypton), the 54th (xenon), and the 86th (radon). Pg 187 • The chemical family called the halogens, composed of elements 9 (fluorine), 17 (chlorine), 35 (bromine), 53 (iodine), and 85 (astatine), is an extremely reactive family. Electron Location • Kernel of an atom is the nucleus and all the electrons but the valence electrons. • The elements are arranged vertically in columns of the periodic table called GROUPS or FAMILIES. • Group # indicates the number of valence electrons. • Because of Periodicity, the elements with the same # of valence electrons are in the same group. • These electrons influence the chemical and physical properties of elements the most. • Electron Configuration shows the location of all the electrons for the atom. Physical and Chemical properties • Because of Periodicity, the elements with the same # of valence electrons are in the same group so they share similar chemical and physical properties. • Chemical properties of matter describe its "potential" to undergo some chemical change or reaction by virtue of its composition. What elements, electrons, and bonding are present to give the potential for chemical change. The result of the change is the formation of a new substance. – Toxicity – Flammability – Reactivity • Electronegativity • Ionization Energy • Physical properties can be observed or measured without changing the composition of matter. They are used to observe and describe matter. – – – – – – – – Atomic radius Density Melting Point Boiling Point Color Solubility Odor Conductivity Transition Metals Lanthanide Series Actinide Series Noble Gases or Inert Gases Halogens Alkaline Earth Metals Alkali Metals Because of periodicity…. • The alkali metals are silvercolored • Soft solids (Fr and Cs are liquids) • The first three are biologically important • low-density metals • react readily with halogens • react readily with water • one valence electron – so they want to lose an electron and achieve a noble configuration (which is?) – form +1 cations 1A NOT Alkali Metals Alkali Metals Alkaline Earth Metals – so they want to lose 2 electrons and achieve a noble configuration (an octet) – form +2 cations 1A 2A Alkaline Earth Metals silvery colored Soft solids Ca and Mg have biological functions low-density metals react readily with halogens react readily with water - though not as rapidly as the alkali metals • Beryllium is an exception: It does not react with water • two valence electron Alkali Metals • • • • • • Transition Metals • They often form colored compounds. • They are often good catalysts – lowers activation energy so rxns are faster – not used up in the rxn – enzymes • They are silvery-blue at room temperature (except copper and gold) - lustrous. • Malleable • They are solids at room temperature (except Hg) • Partly filled d sublevel B group Transition Metals – They can have a variety of different charged cations – 4s fills before 3d (clouds are more apparent and overlapping occurs) • • • • chromium Iron Vanadium Silver (5s fills before 4f) • Good conductors of electricity (Why?) Thus they are transitioning between the filling of their outermost orbitals. Transition Metal Colors Lanthanide • Rare Earth Metals • Silvery-white metals that tarnish when exposed to air • Relatively soft metals • Very reactive • Many rare earth compounds fluoresce strongly under ultraviolet light Lanthanide Series Actinide • All are radioactive • The metals tarnish readily in air • Actinides are very dense metals • Actinides combine directly with most nonmetals Actinide Series Metallic Characteristics • • • • • • • • Conduct electricity and heat Dense Malleable (bendable to form shapes – jewelry) Ductile (able to be drawn out into wires) Lustrous (shiny) Reactive High Melting point Solid at RT Metalloids • Share properties with both metals and nonmetals • Solids • Semi-conductors (between a conductor and an insulator) • Form cations or anions. # of valence electrons varies. • 4 sit on steps and 2 are beneath. 8A Less 3A 4A 5A 6A 7A metallic metalloid Poor Metals More metallic metalloid NonMetals 8A 5A 6A 7A Noble gases 4A Halogens • poor conductors of heat and electricity • in solid form, they are dull and brittle • usually have lower densities than metals • most of the crust, atmosphere and oceans are made up of nonmetals. • Bulk tissues of living organisms are composed almost entirely of nonmetals. • Many nonmetals (hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine) are diatomic, and most of the rest are polyatomic • Prefer to form anions - gaining electrons to achieve an octet. Halogens • This high reactivity is due to their atoms being one electron short of a full outer shell. 8A • Halogens are highly reactive • They form diatomic molecules (F2, Cl2, Br2, I2) • All three states of matter are represented – fluorine and chlorine are gases – bromine is a liquid – iodine and astatine are solids Noble Gases • Both chlorine and bromine are used as disinfectants 7A Halogens – harmful or lethal to biological organisms in sufficient quantities. – Fluorine is the most reactive element in existence – This high reactivity is due to their atoms being one electron short of a full outer shell of electrons. They form -1 anions. Diatomics • hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine • Diatomic elements are nonmetal elements that form a covalent bond between two atoms. • As elements they always travel in pairs of atoms and therefore you must write then as: • H2, N2, O2, F2, Cl2, Br2, I2 Noble Gases or Inert Gases – They have the maximum number of valence electrons their outer shell can hold. Noble Inert Gases • odorless, colorless, monatomic gases. • Lighting (Ne), welding and space technology (processes performed under Ar so that no unwanted chem rxns occur) . • Stable or unreactive 8A PEL / PQN • Probable Location according to the wave mechanical model • Principle Energy Levels (PEL) or Principle Quantum Number (PQN) = n – – – – – – – – – – the total number of orbits around the nucleus Period # = n Max number of e in that PEL, PQN = 2n2 When n = 1, max # of e = 2 When n = 2, max # of e = 8 When n = 3, max # of e = 18 When n = 4, max # of e = When n = 5, max # of e = When n = 6, max # of e = When n = 7, max # of e = Sublevels • Sublevels exist in each orbit (PEL) – Shape of the electron cloud that is created by fast moving electrons. – spdfg – The number of sublevels present in each PEL also = n, so PEL 5 contains 5 sublevels. They are 5s, 5p, 5d, 5f, and 5g Element (Neutral) Galium E shown on the periodic table 2 8 18 3 PQN (4th period) 1 2 3 4 Sublevels 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f Orbitals • Each sublevel has orbitals – – – – – The possible orientations of the shapes around the x,y, and z axis s sublevel has 1 orbital p sublevel has 3 orbitals d sublevel has 5 orbitals f sublevel has 7 orbitals Element (Neutral) Galium E shown on the periodic table 2 8 18 3 PQN 1 2 3 4 Sublevels 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f Orbitals 1 1 3 1 3 5 1 3 5 7 s sublevel orbitals p sublevel orbitals d sublevel orbitals Electrons in Orbitals • Each orbital can contain 2 electrons maximum • 2n2 Element (Neutral) Galium E shown on the periodic table 2 8 18 3 PQN 1 2 3 4 Sublevels 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f Orbitals 1 1 3 1 3 5 1 3 5 7 # of E possible 2 2 6 2 6 10 2 6 10 14 Actual # of E 2 2 6 2 6 10 2 1 E configuration 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1 Determining spdf configuration using THE Table The number of columns present in the block equals the number of possible electrons for that sublevel. f Block location if inserted Really know the location of your valence electrons Using PT to determine detailed electron configuration A short cut…? • Diagonal rule - A guideline explaining the order in which electrons fill the orbital levels. • Pros – Easy to use if given to you • Cons – Memorization required – There are exceptions to this rule when filling the orbitals of heavier elements. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Using Diagonal Rule compared to using PT for Determining detailed electron configuration Use PT for determination of Cr Use diagonal rule for determination of Cr Use PT for determination of Pd Use diagonal rule for determination of Pd 7s 7p 7d 7f 7s 7p 7d 7f 6s 6p 6d 6f 6s 6p 6d 6f 5s 5p 5d 5f 5s 5p 5d 5f 4s 4p 4d 4f 4s 4p 4d 4f 3s 3p 3d 3s 3p 3d 2s 2p 2s 2p 1s 1s Summary • Elements are arranged by increasing atomic # across the periodic table. • Periods are the horizontal rows. The period # = PEL. • Elements are grouped vertically by similar chemical and physical properties. • Groups (or Families) are the vertical columns. The group # = # of valence electrons. Textbook pg 186-191 Practice problems 5.1-5.3 Atomic Radius • Half the distance between the nuclei of two like atoms in a diatomic molecule. • Atom size vs ion size Combine to form O2 Pg 187-188 As you go down a group, Atomic Radius increases As you go across a period, Atomic Radius decreases Why? • As you go down a group the number of PELs increases, more electrons are present to fill these energy levels, so atomic radius increases. • As you go across a period, atomic # increases which means that the # of protons in the nucleus increases, so nuclear charge is increasing and attracting electrons with a greater force. Opposite attract. Ionization Energy • The amount of energy required to remove an outer electron • The more difficult it is to remove an electron, the greater the ionization energy • smaller atoms have greater ionization energy since the valence electrons are closer to the nucleus and more strongly attracted and, therefore, more difficult to remove • • X + energy → X+ + e• First Ionization Energy • X + energy → X+ + e- • Second Ionization Energy (greater than 1st) • X+ + energy → X+2 + e- • Third Ionization Energy (greater than 2nd) • X+2 + energy → X+3 + e- Pg 189 As you go down a group, Ionization Energy decreases Group 2 and 18 2100 ionization energy If the ionization energy is high, that means it takes a lot of energy to remove the outermost electron. If the ionization energy is low, that means it takes only a small amount of energy to remove the outermost electron 1600 1100 600 100 -400 0 10 20 30 40 atomic number Top trend is group 18 Bottom trend is group 2 50 60 As you go across a period, Ionization energy increases Why? • As you go down a group the number of PELs increases. Attraction is less for the electrons furthest from the nucleus so it takes less energy for electrons to be pulled away. • As you go across a period, nuclear charge is increasing and attracting electrons with a greater force. Since that force is increasing, it takes more energy for the electrons to be pulled away. Going towards inert gases which have a full valence shell and are extremely resistant to give up any electrons. Electronegativity • An atom’s affinity for electrons • Arbitrary scale from 04 – 0 is least electronegative – 4 is most electronegative Pg 189-191 Fluorine is the most electronegative 4 Francium is the least electronegative 0.7 Neon and the other noble gases have an Electronegativity of 0 As you go down a group, Electronegativity decreases Electronegativity is a measure of the tendency of an atom to attract electrons. The arbitrary scale of 04 is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0. As you go across a period, Electronegativity increases Why? • As you go down a group the number of PELs increases. Electron attraction to the nucleus is less when they are farther from the nucleus. • As you go across a period, nuclear charge is increasing and thus attracting electrons to a greater extent. • Why do the inert gases have an electronegativity of 0? Chemical Reactivity • Francium is the most reactive metal. • Fluorine is the most reactive non-metal. Chemical Reactivity Electronegativity Ionization Energy Elements on opposite Sides of the Periodic Table are attracted to each other. • Sodium likes to combine with Chlorine – Why? • Atoms become cations due to less ionization energy. – METALS – Fr • Atoms become anions due to high electronegativity. – NONMETALS –F