Periodic Properties of the Elements - Pearson Chemistry PDF

Chemistry: The Central Science Fifteenth Edition Chapter 7 Periodic Properties of the Elements Copyright © 2023 Pearson Education, Inc. All Rights Rese...

Chemistry: The Central Science Fifteenth Edition Chapter 7 Periodic Properties of the Elements Copyright © 2023 Pearson Education, Inc. All Rights Reserved 7.1 Development of the Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Mendeleev and the Periodic Table Table 7.1 Comparison of the Properties of Eka-Silicon Predicted by Mendeleev with the Observed Properties of Germanium Mendeleev’s Predictions for Observed Properties of Germanium Property Eka-Silicon (made in 1871) (discovered in 1886) Atomic weight 72 72.59 Density 5.5 5.35 (g/cm3 ) grams per cubic centimeters Specific heat 0.305 0.309 (J/g-K) Joules per gram kelvin Melting point High 947 (°C) degrees Celsius Color Dark gray Grayish white Formula of oxide G e, O 2 XO2 XO2 GeO Density of oxide grams per cubic centimeters 4.7 4.70 2 (g/cm3 ) Formula of chloride XCl4 GeCl 4 XCl4 GeCl Boiling point of chloride degrees Celsius A little under 100 84 4 (°C) Chemists mostly credit Mendeleev because he also used chemical properties to organize the table and predicted some missing elements and their expected properties, including germanium. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Atomic Number Mendeleev’s table was based on atomic mass. It was the most fundamental property of elements known at the time. About 35 years later, the nuclear atom was discovered by Ernest Rutherford. Henry Moseley developed the concept of atomic number experimentally using X-rays. The number of protons was considered the basis for the periodic property of elements. – Mosley also made it possible to identify “holes” in the periodic table which led to discovery of new elements. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Periodicity Periodicity is the repetitive pattern of a property for elements based on atomic number. The following properties are discussed in this chapter: – Sizes of atoms and ions – Ionization energy – Electron affinity – Some group chemical property trends First, we need to present a fundamental property that led to many of the trends. This is effective nuclear charge. Copyright © 2023 Pearson Education, Inc. All Rights Reserved 7.2 Effective Nuclear Charge (1 of 2) Many atomic properties depend on attractions between valence electrons and the nucleus. Electrons are both attracted to the nucleus and repelled by other electrons. The forces an electron experiences depend on both factors. The nuclear charge is screened. Copyright © 2023 Pearson Education, Inc. All Rights Reserved 7.2 Effective Nuclear Charge (2 of 2) The screen or effective nuclear charge, Zeff , is determined using: Zeff Z  S where Z is the atomic number and S is a screening constant. S is usually the number of core electrons. Effective nuclear charge is a periodic property: – Increases left to right across a period. Copyright © 2023 Pearson Education, Inc. All Rights Reserved 7.3 Sizes of Atoms and Ions The space occupied by electrons is a probability. Need to decide on the “edge”. The nonbonding atomic radius, or van der Waals radius, is half of the shortest distance separating two nuclei during a collision of atoms. The bonding atomic radius, or covalent radius, is half the distance between nuclei in a bond. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Bonding Atomic Radius The bonding atomic radius tends to – decrease from left to right across a period (Zeff  ). – increase from top to bottom of a group (n  ). Copyright © 2023 Pearson Education, Inc. All Rights Reserved Sizes of Ions—Ionic Radii (1 of 2) Determined by interatomic distances in ionic compounds Ionic size depends on – the nuclear charge; – the number of electrons; – the orbitals in which valence electrons reside. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Sizes of Ions—Ionic Radii (2 of 2) Cations are smaller than their parent atoms: – The outermost electron(s) is/are removed and repulsions between electrons are reduced. Anions are larger than their parent atoms: – Electrons are added and repulsions between electrons are increased. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Electron Configurations of Ions Cations: The electrons are lost from the highest energy level (n value). – Li is 1s2 2 s1 Li+ is 1 s2 (loss of a 2s electron). – Fe is 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Fe2 is 1s2 2s2 2p6 3s2 3p6 3d6 (loss of a two 4s electrons). Anions: The electron configurations are filled to ns 2np 6 ; for example, F is 1s 2 2s 2 2 p 6 (gaining one electron in 2p). Copyright © 2023 Pearson Education, Inc. All Rights Reserved Isoelectronic Series and Ion Size In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge. – An isoelectronic series example (Neon—10 electrons) Also note the increasing nuclear charge with decreasing ionic radius as atomic number increases. Increasing nuclear charge 8 protons 9 protons 11 protons 12 protons 13 protons 10 electrons 10 electrons 10 electrons 10 electrons 10 electrons O 2 O 2 minus F F minus NaN a plus Mg2M g 2 plus Al3A l 3 plus 1.26 Å 1.26 Angstroms 1.19 Å 1.19 Angstroms 1.16 Å 1.16 Angstroms 0.86 Å 0.86 Angstroms 0.68 Å 0.68 Angstroms Decreasing ionic radius Copyright © 2023 Pearson Education, Inc. All Rights Reserved 7.4 Ionization Energy and Electron Affinity The ionization energy (I) is the minimum energy required to remove an electron from the ground state of a gaseous atom or ion. – The first ionization energy I1  is the energy required to remove the first electron. – The second ionization energy I2  is the energy required to remove the second electron. The higher the ionization energy, the more difficult it is to remove an electron. You are removing a negative particle from a positive ion. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Ionization Energy More energy is required to remove each successive electron. When all valence electrons have been removed, it takes a great deal more energy to remove the next electron (a core electron). Table 7.2 Successive Values of Ionization Energies, I, for the Elements Sodium through Argon kJ/mol  Element I1 l sub 1 I2 l sub 2 I3 l sub 3 I3 l sub 3 I3 l sub 3 I4l sub 4 I5 l sub 5 Na 496 4562 Blank Blank Blank Blank Blank Mg 738 1451 7733 Blank Blank (inner-shell electrons) Blank Al 578 1817 2745 11,577 Blank Blank Blank Si 786 1577 3232 4356 16,091 Blank Blank P 1012 1907 2914 4964 6274 21,267 S 1000 2252 3357 4556 7004 8496 27,107 Cl 1251 2298 3822 5159 6542 9362 11,018 Ar 1521 2666 3931 5771 7238 8781 11,995 Copyright © 2023 Pearson Education, Inc. All Rights Reserved Periodic Trends for the First Ionization Energy (I1), I sub 1 1) I1 generally increases across a period. 2) I1 generally decreases down a group. 3) The s- and p-block elements show a larger range of values for I1. (The d-block generally increases slowly across the period; the f-block elements show only small variations.) Copyright © 2023 Pearson Education, Inc. All Rights Reserved Irregularities in the General Trend The trend is not followed when the added valence electron in the next element – enters a new sublevel (higher energy sublevel); – is the first electron to pair in one orbital of the sublevel (electron repulsions lower energy). Copyright © 2023 Pearson Education, Inc. All Rights Reserved Factors That Influence Ionization Energy Smaller atoms have higher I values. I values depend on effective nuclear charge and average distance of the electron from the nucleus. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Electron Affinity (EA) Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl(g)  e  Cl (g) EA  349 kJ It is typically exothermic, so, for most elements, it is negative. Copyright © 2023 Pearson Education, Inc. All Rights Reserved General Trend in Electron Affinity Not as established as ionization. Across a period, it generally increases. Three notable exceptions include the following: 1) Group 2A: s sublevel is full. 2) Group 5A: p sublevel is half- full. 3) Group 8A: p sublevel is full. *Note: The electron affinity for many of these elements is positive (X is unstable). Copyright © 2023 Pearson Education, Inc. All Rights Reserved 7.5 Metal, Nonmetals, and Metalloids Copyright © 2023 Pearson Education, Inc. All Rights Reserved Metals Differ from Nonmetals Metals tend to form cations. Nonmetals tend to form anions. Note the special property of hydrogen, i.e., H and H. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Metals Most of the elements in nature are metals. Properties of metals: – Shiny luster – Conduct heat and electricity – Malleable and ductile – Solids at room temperature (except mercury) – Low ionization energies/form cations easily Gold Copyright © 2023 Pearson Education, Inc. All Rights Reserved Metal Chemistry Compounds formed between metals and nonmetals tend to be ionic. Metal oxides tend to be basic and react with acids. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Nonmetals Nonmetals are found on the right hand side of the periodic table. Properties of nonmetals include the following: – Solid, liquid, or gas (depends on element) – Solids are dull, brittle, poor conductors – Large negative electron affinity, so they form anions readily Sulfur Copyright © 2023 Pearson Education, Inc. All Rights Reserved Nonmetal Chemistry Solid carbon dioxide (CO2 , dry ice) in water. Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Comparison of the Properties of Metals and Nonmetals Table 7.3 Characteristic Properties of Metals and Nonmetals Metals Nonmetals Have a shiny luster; various colors, Do not have a luster; various colors although most are silvery Solids are malleable and ductile Solids are usually brittle; some are hard, and some are soft Good conductors of heat and Poor conductors of heat and electricity electricity Most metal oxides are ionic solids that Most nonmetal oxides are molecular are basic substances that form acidic solutions Tend to form cations in aqueous Tend to form anions or oxyanions in solution aqueous solution Copyright © 2023 Pearson Education, Inc. All Rights Reserved Metalloids Metalloids have some characteristics of metals and some of nonmetals. Several metalloids are electrical semiconductors (computer chips). Elemental silicon Copyright © 2023 Pearson Education, Inc. All Rights Reserved Group Trends Elements in a group have similar properties. Trends also exist within groups. Groups compared: – Group 1A: the alkali metals – Group 2A: the alkaline earth metals – Group 6A: the oxygen group – Group 7A: the halogens – Group 8A: the noble gases – Why hydrogen is a nonmetal Copyright © 2023 Pearson Education, Inc. All Rights Reserved 7.6 Group A1: Alkali Metals Alkali metals are soft, metallic solids. They are found only in compounds in nature, not in their elemental forms. Typical metallic properties (luster, conductivity) are seen in them. Sodium Copyright © 2023 Pearson Education, Inc. All Rights Reserved Alkali Metal Properties They have low densities and melting points. They also have low ionization energies. Table 7.4 Some Properties of the Alkali Metals Electron Element Configuration Melting Point (°C) Density (g , degrees Celsius cm3 ) grams per cubic centimeters Atomic Radius (Å) I1(kJ mol) Angstroms I sub 1, kilo Joules per mole Lithium 181 0.53 1.28 520 He 2s1 left bracket H e right bracket 2 s to the first power Sodium 98 0.97 1.66 496 Ne 3s1 left bracket N e right bracket 3 s to the first power Potassium 63 0.86 2.03 419  Ar  4s1 left bracket ay r right bracket 4 s to the first power Rubidium 39 1.53 2.20 403 Kr  5s1 left bracket K r right bracket 5 s to the first power Cesium 28 1.88 2.44 376  Xe 6s1 left bracket X e right bracket 6 s to the first power Copyright © 2023 Pearson Education, Inc. All Rights Reserved Alkali Metal Chemistry Their reactions with water are famously exothermic. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Differences in Alkali Metal Chemistry Lithium reacts with oxygen to make an oxide: 4 Li  O2 2Li2O Oxide ion : O2  Sodium reacts with oxygen to form a peroxide: 2 Na  O2 Na2O2 Peroxide ion : O 22 K, Rb, and Cs also form superoxides: M  O2 MO2 Superoxide ion : O 2  Copyright © 2023 Pearson Education, Inc. All Rights Reserved Flame Tests Qualitative tests for alkali metals include their characteristic colors in flames. These are caused by electronic transitions. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Group 2A: Alkaline Earth Metals (1 of 2) Table 7.5 Some Properties of the Alkaline Earth Metals Electron Atomic I1(kJ mol) Configuratio (°C) (g cm3 ) Radius (Å) I sub 1, kilo Joules per mole Element n Melting Point Density Angstroms , degrees Celsius grams per cubic centimeters Beryllium He 2s 2 left bracket H e right bracket 2 s squared 1287 1.85 0.96 899 Magnesium Ne 3s 2 left bracket N e right bracket 3 s squared 650 1.74 1.41 738 Calcium  Ar  4s 2 left bracket A r right bracket 4 s squared 842 1.55 1.76 590 Strontium Kr  5s 2 left bracket K r right bracket 5 s squared 777 2.63 1.95 549 Barium  Xe 6s 2 left bracket X e right bracket 6 s squared 727 3.51 2.15 503 Alkaline earth metals have higher densities and melting points than alkali metals. Their ionization energies are low, but not as low as those of alkali metals. They readily form +2 cations, losing the 2 valence electrons. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Alkaline Earth Metals (2 of 2) Beryllium does not react with water, and magnesium reacts only with steam, but the other alkaline earth metals react readily with water. Reactivity tends to increase as you go down the group. Electrons are further form the nucleus and more readily react. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Group 6A—Increasing in Metallic Character Down the Group Table 7.6 Some Properties of the Group 6A Elements Electron Element Configuration Melting Point (°C) , degrees Celsius Density Atomic Radius (Å) Angstroms I1(kJ mol) I sub 1, kilo Joules per mole Oxygen left bracket H e right bracket 2 s squared 2 p to the fourth power 1.43 g/L 0.66 1314 He 2s 2 2p 4  218 negative 218 1.43 grams per Liter Sulfur left bracket N e right bracket 3 s squared 3 p to the fourth power 115 1.05 1000 Ne 3s 2 3 p 4 1.96 g/cm3 1.96 grams per cubic centimeter Selenium left bracket A r right bracket 3 d to the tenth power 4 s squared 4 p to the fourth power 221 1.20 941  Ar  3d 10 4s 2 4 p 4 4.82 g/cm3 4.82 grams per cubic centimeter Tellurium left bracket K r right bracket 4 d to the tenth power 5 s squared 5 p to the fourth power 450 1.38 869 Kr  4d 10 5s 2 5 p 4 6.24 g/cm3 6.24 grams per cubic centimeter Polonium left bracket X e right bracket 4 f to the fourteenth power 5 d to the tenth power 6 s squared 6 p to the fourth power 254 1.40 812  Xe 4f 5d 6s 6 p 9.20 g/cm3 14 10 2 4 9.20 grams per cubic centimeter Oxygen, sulfur, and selenium are nonmetals. Tellurium is a metalloid. The radioactive polonium is a metal. Trend: Oxygen is more likely to form  2 anion; polonium is most likely to have a positive charge. Copyright © 2023 Pearson Education, Inc. All Rights Reserved 7.7 Select Nonmetals: Hydrogen​ Electron configuration of 1s1 suggests it should be a metal, yet IE 1312 kJ does not fit.​ Occurs as a colorless diatomic gas, H2. In the presence of water readily forms H. When reacting with metals, can gain electrons to form hydride anions, H. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Oxygen Occurs as the colorless diatomic gas molecule, O2. Oxygen can exist as two different elemental forms also known as allotropes: – Oxygen, O2 – colorless and odorless gas – Ozone, O3 – pale blue, poisonous gas with sharp odor Ions are oxides (O2 ), peroxides (O22 ) and superoxides (O2  ). Copyright © 2023 Pearson Education, Inc. All Rights Reserved Group 7A—Halogens Table 7.7 Some Properties of the Halogens Electron Element Configuration Melting Point (°C) Density , degrees Celsius Atomic Radius (Å) I (kJ mol) Angstroms 1 I sub 1, kilo Joules per mole Fluorine 0.57 1681 left bracket H e right bracket 2 s squared 2 p to the fifth power He 2s 2 2p5  220 negative 220 1.69 g/L 1.69 grams per liter Chlorine 1.02 1251 left bracket N e right bracket 3 s squared 3 p to the fifth power Ne 3s 2 3 p5  102 negative 102 3.12 g/L 3.12 grams per liter Bromine 1.20 1140 left bracket A r right bracket 4 s squared 3 d to the tenth power 4 p to the fifth power  Ar  4s 2 3d 10 4 p5  7.3 3.12 g/cm3 negative 7.3 3.12 grams per cubic centimeter Iodine 114 1.39 1008 left bracket K r right bracket 5 s squared 4 d to the tenth power 5 p to the fifth power Kr  5s 2 4d 10 5 p5 4.94 g/cm3 4.94 grams per cubic centimeter The halogens are typical nonmetals. They have highly negative electron affinities, so they exist as anions in nature. They react directly with metals to form metal halides. Copyright © 2023 Pearson Education, Inc. All Rights Reserved  Group 8A—Noble Gases Table 7.8 Some Properties of the Noble Gases Electron Element Configuration Boiling Point (K) Density (g/L) Atomic Radius* (Å) I (kJ mol) , grams per liter Angstroms 1 I sub 1, kilo Joules per mole Helium 4.2 0.18 0.28 2372 1 s squared 1s 2 Neon 27.1 0.90 0.58 2081 left bracket H e right bracket 2 s squared 2 p to the sixth power He 2s 2 2p6 Argon 87.3 1.78 1.06 1521 left bracket N e right bracket 3 s squared 3 p to the sixth power Ne 3s 2 3 p6 Krypton 120 3.75 1.16 1351 left bracket ay r right bracket 4 s squared 3 d to the tenth power 4 p to the sixth power Xenon  Ar  4s 2 3d 10 4 p6 left bracket K r right bracket 5 s squared 4 d to the tenth power 5 p to the sixth power 165 5.90 1.40 1170 Radon Kr  5s 2 4d 5 p 10 6 left bracket X e right bracket 6 s squared 4 f to the fourteenth power 5 d to the tenth power 6 p to the sixth power 211 9.73 1.50 1037  Xe 6s 2 4f 14 5d 10 6 p6 *Only the heaviest of the noble-gas elements form chemical compounds. Thus, the atomic radii for the lighter noble-gas elements are estimated values. Have very large ionization energies—full electron shells (don’t form cations). Their electron affinities are positive (can’t form stable anions). Therefore, they are relatively unreactive. All are monatomic gases. Copyright © 2023 Pearson Education, Inc. All Rights Reserved Copyright This work is protected by United States copyright laws and is provided solely for the use of instructors in teaching their courses and assessing student learning. Dissemination or sale of any part of this work (including on the World Wide Web) will destroy the integrity of the work and is not permitted. The work and materials from it should never be made available to students except by instructors using the accompanying text in their classes. All recipients of this work are expected to abide by these restrictions and to honor the intended pedagogical purposes and the needs of other instructors who rely on these materials. 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