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CHEMICAL PERIODICITY ANTOINE LAVOISIER (1790s) “The father of modern chemistry The first to recognize the true elements JOHANN DOBEREINER (1817) Found that several groups of three elements have similar properties and called it TRIADS Second element has an atomic mass halfway between 1st and 3rd elements STANISLAO CANNIZZARO (1860) Presented an accurate method for measuring relative masses of atoms at the First International Congress of Chemists in Karlsruhe, Germany This initiated a search for the relationships between atomic mass and properties of elements Clearly distinguished between atoms and molecules JOHN NEWLANDS (1863-1864) Arranged the elements in order of their increasing atomic masses Elements were arranged in sequential order, assigning an ordinal number to each Noted that there is repetition of similar properties every eighth element Placed seven elements in each group and called it OCTAVES LAW OF OCTAVES: _______________________________________________ Did not work for all elements DMITRI MENDELEEV (1869) Listed 60+ elements in several vertical columns in order of their increasing atomic mass Noticed a regular recurrence of their physical and chemical properties ____________________ in the table because there were no known elements with the appropriate properties at that time 1871 - Predicted the physical and chemical properties of 3 missing elements _________________________ German chemist Working on element organization at the same time as Mendeleev Published his work a year later than Mendeleev WILLIAM RAMSEY 1868-1900 Discovery of noble gases Proposed a new group in the Periodic Table to accommodate the noble gases discovered Noble gases were hard to detect due to __________________________________ HENRY MOSELEY (1911-1913) Mendeleev had a few elements that did not fit Determined the nuclear charge and atomic number of elements by analyzing the spectra of 38 metals Rearranged the elements in the table in the order of their __________________________ HARRY D. HUBBARD (1924) Modernized Mendeleev's periodic table His first work was published in 1924 known as the "Periodic Chart of the Atoms" POshikiri / Chemistry 1 of 9 GLENN SEABORG (1940) Discovered 10 Transuranium elements Reorganized the table in 1944 by placing the lanthanide/actinide series at ________________ of the table MODERN PERIODIC TABLE PERIODIC TABLE Arrangement of elements in order of their increasing atomic numbers so that elements with similar properties fall in the same column or group PERIODIC LAW The physical and chemical properties of the elements are periodic functions of their ______________________ PERIOD or SERIES Horizontal rows of the periodic table 7 periods Denote # of _________________________ GROUP or FAMILY Vertical columns of the periodic table Denote # of _________________________ POshikiri / Chemistry 2 of 9 POshikiri / Chemistry 3 of 9 ELECTRON CONFIGURATION & PT MAIN GROUP ELEMENTS s- and p-block elements 1A to 8A (1,2,,13,14,15,16,17,18) REPRESENTATIVE ELEMENTS Group A elements in the periodic table 1A, 2A, 3A, 4A, 5A, 6A, 7A s and p sublevels are partially filled s-Block: Outermost electrons are in the s-sublevel Alkali Metals – very active metals Alkaline Earth Metals – also active Hydrogen and Helium p-Block Groups 13 – 18 Electrons add to a p-sublevel only after the s- sublevel in the same energy level is filled Boron, carbon, nitrogen, oxygen, halogens,noble gas families HALOGENS – salt formers, very reactive NM METALLOIDS – brittle solids with some properties of both metals and nonmetals NOBLE GASES – inert due to filled s and p orbitals d-Block Group B elements in the periodic table/Groups 3-12 Transition Elements – d-block elements with typical metallic properties Form colored ions f-Block Lanthanides: 4f sublevel is being filled Actinides: 5f sublevel is being filled POshikiri / Chemistry 4 of 9 PERIODIC PROPERTIES Shielding Effect Atomic Radius Ionization Energy Electron Affinity Ionic Radii Electronegativity ATOMIC FORCES NUCLEAR PULL Theoretical force equal to Z Z = # of protons ELECTRON REPULSION SHIELDING Inner shells of e- “shield” the outer shell e- from the nucleus’ pull EFFECTIVE NUCLEAR CHARGE EFFECTIVE NUCLEAR CHARGE (Zeff) is the “_______________________________________________. Outer electrons feel less than full strength of nucleus because of electrons between them Z (shielding effect) = Zeff Example: Na (Z = 11) Filled 1s2s2p effectively blocks out 10p+ Valence e- only feels 1 p+ Zeff = 1 Within a period, as Zeff ______________, radius ________________ SHIELDING EFFECT Decrease in the attraction between outer electrons and the nucleus due to the presence of other electrons between them PERIOD TREND Constant within a period because the _________________________________ GROUP TREND Increases from top to bottom because of the increase in the main energy levels ATOMIC RADIUS _____________________ between the nuclei of identical atoms that are bonded together GROUP TREND Atomic radii of the main group elements increase down a group This is due to _________________ in the number of main energy levels Because electrons are added further from the nucleus, there is _____________________. Size of the atom is generally determined by the number of energy levels POshikiri / Chemistry 5 of 9 PERIOD TREND Decreasing trend across a period Across a period, there is an increase of protons in the nucleus, and also an increase of electrons in the same main energy level Size ____________________________________________. Each added electron feels a greater and greater + charge. The increase in nuclear charge increases the attraction of the nucleus for the outermost electrons, pulling it tightly towards the nucleus, causing the size of the atom to decrease IONIZATION ENERGY Energy required to _______________________________ Or how tightly an e- is held by atom ION – atom or group of bonded atoms that has a positive or negative charge IONIZATION – process that results in the formation of ion First ionization energy (IE1) Minimum amount of energy needed to remove the most loosely held e- from an atom Ex: Ca + 590 kJ → Ca1+ + eSecond ionization energy (IE2) Amount of energy needed to remove the second eEx: Ca1+ + 1145 kJ → Ca2+ + eIE2 is always greater than IE1 ____________________________ to remove e- from +ion than a neutral atom FACTORS AFFECTING IONIZATION ENERGY NUCLEAR CHARGE The larger the nuclear charge, the greater the ionization energy SHIELDING EFFECT The greater the shielding effect, the less the ionization energy RADIUS The greater the distance between the nucleus and outer electrons, the less the ionization energy SUBLEVEL An electron from a full or half-full sublevel requires additional energy to be removed GROUP TREND ___________________ down a group Size of the atom increases as we go down The outermost electron is farther away from the nucleus, thus there is decreased attraction between nucleus and outermost electrons Outermost electrons are easily removed and therefore have a lower ionization energy PERIOD TREND Ionization energies of main group elements ___________________________ Nuclear charge increases within a period, which strongly attracts the added electrons in the same energy level The atoms become smaller; outermost electrons are held tightly to the nucleus, making it harder to remove Slight decrease between IIA and IIIA due to p sublevels having a higher energy so electrons are easier to remove POshikiri / Chemistry 6 of 9 Slight decrease between VA and VIA due to paired electrons in the p sublevel. Paired electrons have greater repulsive forces so they are easier to remove Noble gases have highest IE (don’t want to let go of electrons) IE helps predict whether an element is likely to form an ___________________________________. Metals tend to have ______ IE and form cations, while nonmetals tend to have _____ IE and form anions – thus forming ionic bonds Elements with intermediate IE form molecular/covalent compounds by sharing electrons ELECTRON AFFINITY Energy change that occurs when a gas phase atom gains an electron to form a gas phase anion. A(g) + e- ---> A-(g) E.A. = ∆E Or atom’s attraction for additional electrons PERIOD TREND _____________ as we move left to right across a period The atoms become smaller and the nuclear charge increases; outermost electrons are held tightly to the nucleus Increasingly negative left to right across a period GROUP TREND Electron affinity ________________ as we move down a group because of the increasing atomic size, which decreases attraction between the nucleus and the outermost electrons Increasingly positive down a group so electrons add with greater difficulty Metals have low electron affinities Nonmetals have high electron affinities Exceptions with IIA, VA, and VIIIA Why? Increased stability IIA – full s VA – ½ full p VIIIA – full s and p Decreased tendency to gain/lose eIONIC RADII REMEMBER THAT ALL ATOMS: Want to have a noble gas configuration Achieve maximum stability with lowest energy configuration Gain/lose/share e- to make it happen Formation of Ions – electrons are removed from the outermost main energy level CATION (POSITIVE ION) Smaller than neutral atom because of the loss of outer shell electrons Na: 1s22s22p63s1 Na+: 1s22s22p6 – same as [Ne] When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 2+ 0 6 6 Fe : [Ar]4s 3d or [Ar]3d Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5 POshikiri / Chemistry 7 of 9 ANION (NEGATIVE ION) Always larger than neutral atoms because nuclear attraction is spread over additional electrons Increased electron-electron repulsion causes the species to become larger F-: 1s22s22p6 or [Ne] PERIOD TREND ______________ in the size of positive ions from left to right across a period From Group 15, negative ions (which are much larger in size), gradually decreases from left to right For d-block and f-block elements, electrons are removed from s-sublevel before d- or f- electrons, resulting in formation of different kinds of ions for same element GROUP TREND ___________ in ionic radii for both anions and cations as you go down each group Outer electrons in both cations and anions are in higher energy levels, so there is a gradual increase of ionic radii ELECTRONEGATIVITY Measure of the ability of an atom in a chemical compound to attract electrons Electronegativity values help predict the type of bonding that can exist between atoms in compounds PERIOD TRENDS ___________________ from left to right due to increase in nuclear pull Metallic elements - very low EN Fr & Cs most reactive metals Nonmetallic elements - have high EN F - highest EN = most reactive NM GROUP TRENDS _________________ from top to bottom within a group due to increased shielding effect VALENCE ELECTRONS Electrons in the _______________________________, which are available to be lost, gained or shared in the formation of chemical compounds OXIDATION NUMBERS The ___________________________________ in a molecule if electrons were transferred completely in the direction indicated by the _____________________________ Metals are found on the left and center of the periodic table Their atoms tend to lose electrons and thus have positive oxidation numbers Transition metals can form as many as 4 different cations because of their complex electron arrangement Nonmetals are on the right side of the periodic table Their atoms tend to gain electrons and thus have ___________________________ POshikiri / Chemistry 8 of 9 OXIDES Metallic oxides form ________________________ Nonmetallic oxides form _____________________ ______________ – a substance that has the properties of an acid and a base METALLIC CHARACTER POshikiri / Chemistry 9 of 9