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CH – 5 THE PERIODIC LAW I. HISTORY OF THE PERIODIC TABLE A. Antoine Lavoisier (1790s) 1. Compiled list of 23 known elements B. Johann Dobereiner (1817) 1. Found that several groups of three elements have similar properties and called it triads. 2. Second element has an atomic mass halfway between 1st and 3rd elements EXAMPLES OF DOBEREINER'S TRIADS TRIAD 1 Name TRIAD 2 TRIAD 3 Atomic Atomic Atomic Name Name Mass Mass Mass First Element Calcium 40.1 Chlorine 35.5 Sulfur Third Element Barium 137.3 Iodine 126.9 Tellurium 127.6 Average 88.7 Average 81.2 Average 79.9 Second Element Strontium 87.6 Bromine 79.9 Selenium 32.1 79.0 C. Stanislao Cannizzaro (1860) – presented an accurate method for measuring relative masses of atoms at the First International Congress of Chemists in Karlsruhe, Germany 1. This initiated a search for relationships between atomic mass and properties of elements PL Oshikiri /04 1 D. Emile Beguyer de Chancourtois – Telluric Screw 1. Positioned the elements according to increasing atomic weight along a spiral inscribed on the surface of a cylinder and inclined at 45° from the base. E. John Newlands (1863-1864) 1. Arranged the elements in order of their increasing atomic masses 2. Noted that there is repetition of similar properties every eighth element 3. Placed seven elements in each group, and called it octaves 4. Law of Octaves – the same properties are repeated every eighth element 5. Did not work for all elements NEWLAND'S LAW OF OCTAVES 1 Li Na K 2 Be Mg 3 B Al 4 C Si 5 N P 6 O S 7 F Cl F.Dmitri Mendeleev (1869) 1. Listed 60+ elements in several vertical columns in order of their increasing atomic mass 2. Noticed a regular recurrence of their physical and chemical properties PL Oshikiri /04 2 3. Left blank spaces in the table because there were no known elements with the appropriate properties at that time 4.1871 – Predicted the physical and chemical properties of 3 missing elements 5. Credited with the Periodic Law G. Lothar Meyer (1870) 1. German chemist, working on element organization at the same time as Mendeleev; published his work a year later H. William Ramsey (1868-1900) 1. Discovery of noble gases 2. Proposed a new group in the periodic table to accommodate the noble gases discovered 3. Noble gases were hard to detect due to lack of chemical reactivity I. Henry Moseley and the Periodic Law 1. Mendeleev had a few elements that did not fit 2. Determined the nuclear charge and atomic number of elements by analyzing the spectra of 38 metals 3. Rearranged the elements in the table in the order of their increasing atomic numbers J. Lanthanides & Actinides 1. Lanthanides were separated and identified in the early 1900s a. Hard to isolate because they have very similar physical and chemical properties 2. Actinides series elements were discovered PL Oshikiri /04 3 K. Harry D. Hubbard (1924) 1. Modernized Mendeleev’s periodic table L. Glenn Seaborg (1940) 1. Discovered 10 Transuranium elements 2. Reorganized the table in 1944 by placing the lanthanide/actinide series at the bottom of the table M. Modern Periodic Table 1. Periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic numbers 2. Periodic table – arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group 3. Period or Series – horizontal rows of the periodic table (seven periods) a. Denote number of principal energy levels 4. Groups or Families – vertical groups a. Denote # of valence electrons 5. Noble Gases, Lanthanides, Actinides II. ELECTRON CONFIGURATION AND THE PERIODIC TABLE A. Sublevel Blocks of the Periodic Table 1. a. Main-group elements – s- and p-block elements b. Representative elements – Group A elements 1A – 7A groups; s and p sublevels are partially filled 2. s-block – outermost electrons are in the s sublevel a. Group 1 – Alkali Metals b. Group 2 – Alkaline Earth Metals PL Oshikiri /04 4 c. Hydrogen and Helium 2. p-block – Groups 13-18 a. Electrons add to a p level only after the s sublevel in the same energy level is filled b. Boron, carbon, nitrogen, oxygen & noble gas families c. Halogens (group 17) – salt formers; very reactive nonmetals d. Metalloids – brittle solids with some properties of both metals and nonmetals e. Noble Gases – inert due to filled s and p orbitals 3. d-block – Groups 3-12 (Group B elements) a. Transition elements – d-block elements with typical metallic properties b. Form colored ions 4. f-block – lanthanides and actinides a. Lanthanides – 4f sublevel is being filled b. Actinides – 5f sublevel is being filled III. ELECTRON CONFIGURATION AND PERIODIC PROPERTIES A. Atomic Forces 1. Nuclear Pull – theoretical force equal to Z (atomic #) 2. Electron Repulsion – like charges repel 3. Shielding a. Inner shells of electrons “shield” the outer shell electrons from the nucleus” pull 4. Z-effective PL Oshikiri /04 5 a. Outer electrons feel less than full strength of nucleus because of electrons between them B. Shielding Effect – decrease in the attraction between outer electrons and the nucleus due to the presence of other electrons between them 1. Period Trend – Constant within a period because the main energy level remains the same 2. Group Trend – Increases from top to bottom because of the increase in the main energy level C. Atomic Radius – defined as one-half the distance between the nuclei of identical atoms that are bonded together 1. Period Trends – Decreasing trend across a period. a. Size become smaller across a period due to increase in the number of protons in the nucleus but shielding effect is constant for the electrons being added to same energy level 2. Group trends – Atomic radii of the main-group elements increase down a group a. This is due to increase in the number of main energy levels, thus decreasing the attraction between the nucleus and outermost electrons D. DIATOMIC ELEMENTS – always exist as pairs Br2, I2, N2, Cl2, H2, O2, F2 E. Ionization Energy – energy required to remove one electron from a neutral atom of an element; describes how tightly an electron is held by atom PL Oshikiri /04 6 1. Ion – atom or group of bonded atoms that has a positive or negative charge 2. Ionization – process that result in the formation of an ion 3. First Ionization Energy (IE1) – energy needed to remove the most loosely held electrons from an atom Ca + 590 kJ -- Ca+1 + e4. Second Ionization Enrgy (IE2) – amount of energy needed to remove the second electron Ca+1 + 1145 kJ -- Ca+2 + ea. IE2 is always greater than IE1 ; More difficult to remove e- from +ion than a neutral atom 5. Factors Affecting Ionization Energy a. Nuclear charge – the larger the nuclear charge, the greater the ionization energy b.Shielding effect – the greater the shielding effect, the less the ionization energy c. Radius – the greater the distance between the nucleus and outer electrons, the less the ionization energy d. Sublevel – an electron from a full or half-full sublevel requires additional energy to be removed 6. a. IE helps predict whether an element is likely to form an ionic or covalent compound b. Metals tend to have low IE and form cations, while nonmetals tend to have high IE and form anions – thus causing ionic bonds PL Oshikiri /04 7 c. Elements with intermediate IE form molecular or covalent compounds by sharing electrons PL Oshikiri /04 8 7. Period trends – ionization energies of main group elements increase across each period a. Slight increase between IIA and IIIA due to p-sublevels having a higher energy so electrons are easier to remove b. Slight decrease between VA and VIA due to paired electrons in the p-sublevel. Paired e- have greater repulsive forces so are easier to remove. c. Nuclear charge increases within a period, which more strongly attracts additional electrons in the same energy level 8. Group trends – among the main group elements, ionization energies generally decrease down the groups. a. The size of the atom increases as we go down; thus the outermost electron is farther from the nucleus with an increased shielding effect b. It should be more easily removed and therefore have lower ionization energy F. Electron Affinity – energy change that occurs when an electron is acquired by a neutral atom; atom’s attraction for additional electron 1. Periodic trends – electron affinity generally increases as we move from left to right across a period because atoms become smaller and the nuclear charge increases a. Increasingly negative left to right across a period 2. Group trends – electron affinity decreases as we move down a group because of the increasing atomic size, which decreases attraction between the nucleus and the outermost electrons a. Increasingly positive down a group so electrons add with greater difficulty PL Oshikiri /04 9 3. Metals have low EA; Nonmetals have high EA G. Ionic Radii 1. a. All atoms want to have a noble gas configuration b. Achieve maximum stability with lowest energy configuration c. Gain/lose/share electrons to make it happen 2. Cation – positive ion; smaller than neutral atom because of the loss of outer shell electrons 3. Anion – negative ion a. Always larger than neutral atoms because the electrons are not drawn as strongly as they were before the addition of extra electrons b. More repulsion among electrons 4. Period Trends – left to right across a period, there is a decrease in the size of the positive ions a. From Group 15, negative ions (which are much larger in size) gradually decrease in size as you move right b. For d-block and f-block elements, first electrons to be removed are s-electrons before d- or f- electrons, resulting in formation of different kinds of ions for the same element 5. Group Trends – there is an increase in ionic radii with both anions and cations as you go down each group a. Outer electrons in both cations and anions are in higher energy levels from top to bottom of the periodic table, so there is also a gradual increase of ionic radii H. Valence Electrons – electrons in the highest energy level of the atom, which are available to be lost, gained, or shared in the formation of chemical compounds PL Oshikiri /04 10 PL Oshikiri /04 11 I. Oxidation Numbers – indicates the number of electrons that would be lost or gained by the atom 1. Metals are found on the left and center of periodic table a. Tend to lose electrons, thus they have positive oxidation numbers 2. Nonmetals are on the right side of the periodic table a. Tend to gain electrons, thus they have negative oxidation numbers J. Electronegativity – a measure of the ability of an atom in a chemical compound to attract electrons 1. Electronegativity values help predict the type of bonding that can exist between atoms in compounds. 2. Period Trends – Increases from left to right due to increase in nuclear pull a. Metallic elements at the left have very low electronegativities (Fr & Cs most reactive metals) b. Non-metallic elements at the right have high electronegativities (F – most reactive nonmetal) 3. Group Trends – From top to bottom within a group, electronegativity generally decreases. K. Oxides 1. Groups 1 & 2 oxides are generally basic 2. Nonmetallic oxides are generally acidic L. Periodic Properties of d- and f- Block Elements 1. Atomic Radii 2. Ionization Energy 3. Ion Formation and Ionic Radii PL Oshikiri /04 12 4. Electronegativity PL Oshikiri /04 13