I. Development of the Periodic Table
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(2) I. Development of the Periodic Table.
(3) A. Development of the Table 1. J. Newlands (1864) • arranged elements by atomic mass • noticed properties 'repeated' every 8 elements • law of octaves. Newland's table of the elements. At the time, his 'Law of octaves' was ridiculed by his contemporaries. The Society of Chemists did not accept his work for publication..
(4) A. Development of the Table 2. L. Meyer (1869) • demonstrated a connection between atomic mass and properties of elements. Julius Lothar Meyer and his not-so-famous periodic table..
(5) A. Development of the Table 3. D. Mendeleev (1869) • arranged elements by increasing atomic mass • grouped elements with similar properties • left blank spaces for undiscovered elements • predicted the properties of 'missing' elements.
(6) A. Development of the Table 4. H. Moseley (1913) • discovered each element has a unique number of protons called the atomic number • arranged elements by atomic number • Periodic Law - when elements are arranged by increasing atomic number there is a periodic repetition of properties.
(7) B. The Modern Table • arranged by increasing atomic number.
(8) B. The Modern Table • Groups (Families) - vertical columns that have similar chemical and physical properties.
(9) B. The Modern Table • Periods - horizontal rows.
(10) B. The Modern Table • Representative Elements • s & p-blocks • Groups 1, 2, 13-18.
(11) 1. Metals • left of stair-step.
(12) 1. Metals • good Conductors of heat and electricity.
(13) 1. Metals • high Luster when polished.
(14) 1. Metals • Ductile - can be drawn into wires.
(15) 1. Metals • Malleable - can be pounded/rolled into thin sheets. 22k gold leaf applied with an ox hair brush during the process of gilding. A gold nugget of 5 mm in diameter (bottom) can be expanded through hammering into a gold foil of about 0.5 square meter. Toi museum, Japan..
(16) 1. Metals • most are solid at room T.
(17) 1. Metals • H behaves like a metal under very high pressure Gas giant planets such as Jupiter and Saturn may contain large amounts of metallic hydrogen. (depicted in grey).
(18) 1. Metals a) Alkali Metals • Group 1 (except H) • 1 valence e• extremely reactive • usually exist as compounds. Lithium deuteride was used as fuel in the Castle Bravo nuclear device..
(19) 1. Metals b) Alkali Earth Metals • Group 2 • 2 valence e• very reactive Magnesium firestarter (in left hand), used with a pocket knife and flint to create sparks that ignite the shavings..
(20) 1. Metals c) Transition Metals • Groups 3-12 • d-block.
(21) 1. Metals d) Inner Transition Metals • f-block • Lanthanide & Actinide Series.
(22) 2. Nonmetals • right of stair-step.
(23) 2. Nonmetals • generally do not have luster Carbon. Sulfur.
(24) 2. Nonmetals • poor conductors of heat and electricity.
(25) 2. Nonmetals • many are gases at room T.
(26) 2. Nonmetals • solids tend to be Brittle.
(27) 2. Nonmetals • H behaves like a nonmetal at low pressure. The Space Shuttle Main Engine burns hydrogen with oxygen, producing a nearly invisible flame at full thrust..
(28) 2. Nonmetals a) Halogens • Group 17 • 7 valence e• extremely reactive • usually exist as compounds Chlorine. Bromine. Iodine.
(29) 2. Nonmetals b) Noble Gases • Group 18 • 8 valence e- (He has 2) • un-reactive • full outer E level • also called inert gases because they do not participate in many chemical rxns.
(30) 3. Metalloids • touching stair-step.
(31) 3. Metalloids • share properties of metals and nonmetals.
(32) 3. Metalloids • important in electronics and circuit boards.
(33) 3. Metalloids • B, Si, Ge, As, Sb, Te, Po, At.
(34) Silicon Boron. Antimony. Germanium. Arsenic.
(35) II. Classification of Elements.
(36) A. Organizing by Electron Configuration • e- configuration determines properties. Valence Electrons in each Group.
(37) A. Organizing by Electron Configuration 1. Valence Electrons • e- in highest E level • atoms in the same group have the same # of valence e- and similar chemical properties. The Halogen Group has 7 valence electrons..
(38) A. Organizing by Electron Configuration 2. Valence Electrons and Period • the period (row) indicates the number of E levels • Ex: K (row 4) has some e- in the 4th E level.
(39) A. Organizing by Electron Configuration 3. Valence Electrons and the Representative Elements • the # of valence e- can be predicted for each group • * except He.
(40) B. s-, p-, d-, and f-block Elements • the periodic table is divided into 4 distinct sections.
(41) B. s, p, d, and f-block Elements 1. s-block • Gr 1 Li = 1s22s1 • Gr 2 Be = 1s22s2.
(42) B. s, p, d, and f-block Elements 1. s-block • Gr 1 • Gr 2.
(43) B. s-, p-, d-, and f-block Elements 2. p-block • Gr 13 B = 1s22s22p1 • Gr 14 C = 1s22s22p2 • Gr 15 N = 1s22s22p3 • Gr 16 O = 1s22s22p4 • Gr 17 F = 1s22s22p5 • Gr 18* Ne = 1s22s22p6. (*He = 1s2).
(44) B. s-, p-, d-, and f-block Elements 2. p-block • Gr 13 • Gr 14 • Gr 15 • Gr 16 • Gr 17 • Gr 18*. (*except He).
(45) B. s-, p-, d-, and f-block Elements 3. d-block (TM) • Groups 3 - 12 • ‘B’ Group Elements • outermost 's' sublevel filled • 'd' sublevel contain e- (n - 1) • Ex: Ti = [Ar]4s23d2.
(46) B. s-, p-, d-, and f-block Elements 3. d-block (TM) • Groups 3 - 12 • ‘B’ Group Elements • outermost 's' sublevel filled • 'd' sublevel contains e-.
(47) B. s, p, d, and f-block Elements 4. f-block (ITM) • outermost 's' sublevel (partially) filled • 'f' sublevel contains e- (n - 2).
(48) B. s, p, d, and f-block Elements 4. f-block (ITM) • outermost 's' sublevel (partially) filled • 'f' sublevel contains e-.
(49) III. Periodic Trends.
(50) A. Atomic Radius • distance from the center of the nucleus to the outer edge of the e- cloud • 90% probability of finding e- inside the cloud • atoms do not have clearly defined boundaries • atoms can be thought of as 'spheres' • usually measured in picometers (10-12).
(51) A. Atomic Radius 1. Periodic Trends • atomic radius decreases as you go from L R • as you go across a period the principle E level (n) remains the same • the attractive electrical force between the nucleus and e- cloud increases as you move from L R. 191 pm. 160 pm. 130 pm. 118 pm. 110 pm. 102 pm. 95 pm.
(52) A. Atomic Radius 1. Periodic Trends • atomic radius decreases as you go from L R • the attractive electrical force between the nucleus and e- cloud increases as you move from L R. 191 pm. 160 pm. 130 pm. 118 pm. 110 pm. 102 pm. 95 pm.
(53) A. Atomic Radius 2. Group Trends • atomic radius increases as you move down a group • as you move down, e- are added to higher principle E levels (n) • shielding of the nucleus by eincreases as additional orbits are filled. 157 pm. 191 pm. 235 pm. 250 pm. 272 pm.
(54) A. Atomic Radius 2. Group Trends • atomic radius increases as you move down a group • shielding of the nucleus by eincreases as additional orbits are filled. 157 pm. 191 pm. 235 pm. 250 pm. 272 pm.
(55) A. Atomic Radius.
(56) A. Atomic Radius 2. Group Trends • Shielding Effect - describes the decrease in attraction between an e- and the nucleus in any atom with more than one e- shell.
(57) B. Ionic Radius • Ion - an atom (or group of atoms) that has a (+) or (-) charge.
(58) B. Ionic Radius • Ion - an atom (or group of atoms) that has a (+) or (-) charge.
(59) B. Ionic Radius 1. Metals • tend to form (+) ions (cations) • lose e- to obtain a full outer shell • (+) ions are always smaller that the neutral atom • lose e- results in smaller size.
(60) B. Ionic Radius 2. Nonmetals • tend to form (-) ions (anions) • gain e- to obtain a full outer shell • (-) ions are always larger that the neutral atom • gain of e- results in larger size.
(61) B. Ionic Radius 3. Periodic Trends • going from L R there is a gradual decrease in the size of (+) ions • beginning with Group 15 the (-) ions become much larger, then gradually decrease in size from L R.
(62) B. Ionic Radius 4. Group Trends • ionic radius increases as you move down.
(63)
(64) C. Ionization Energy (IE) • E required to remove efrom a gaseous atom • removing 1 e- results in a (+) ion with a charge of +1 • the E needed to remove the 1st e- is not the same as the 2nd, 3rd, etc.
(65) C. Ionization Energy (IE) 1. Metals • low IE • likely to form cations.
(66) C. Ionization Energy (IE) 2. Nonmetals • high IE • likely to form anions.
(67) C. Ionization Energy (IE) 3. Periodic Trends • the first IE increases as you go from L R • attractive electrical force between the nucleus (+) and the cloud (-) increases but the shielding effect is constant.
(68) C. Ionization Energy (IE) 4. Group Trends • the first IE decreases as you move down a group • the shielding effect makes it easier to remove the outermost e- from elements near the bottom of the chart.
(69) C. Ionization Energy (IE) 5. Octet Rule - atoms gain, lose, or share e- in order to acquire a full set of 8 valence e• all noble gases have an octet (He = duplet) • useful for predicting the type of ion.
(70) C. Ionization Energy (IE) 5. Octet Rule - atoms gain, lose, or share e- in order to acquire a full set of 8 valence e• all noble gases have an octet (He = duplet) • useful for predicting the type of ion • Gr 1 - 13 tend to lose e- (cations).
(71) C. Ionization Energy (IE) 5. Octet Rule - atoms gain, lose, or share e- in order to acquire a full set of 8 valence e• all noble gases have an octet (He = duplet) • useful for predicting the type of ion • Gr 1 - 13 tend to lose e- (cations) • Gr 15 - 17 tend to gain e- (anions).
(72) D. Electronegativity (χ) • the tendency of an atom to attract e- from another atom when forming bonds • have been calculated and are expressed as numbers between 0.70 (weak) 3.98 (strong).
(73) D. Electronegativity (χ) • noble gases are omitted because they do not tend to form compounds.
(74) D. Electronegativity (χ) • metals tend to have low χ and lose e- when forming bonds • nonmetals tend to have high χ and gain e- when forming bonds.
(75) D. Electronegativity (χ) 1. Periodic Trend • χ increases as you go from L. R.
(76) D. Electronegativity (χ) 2. Group Trend • χ decreases as you go down a group.
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