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- 1. Slide 1 of 35 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9 PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY TENTH EDITION GENERAL CHEMISTRY Principles and Modern Applications PETRUCCI HERRING MADURA BISSONNETTE The Periodic Table and Some Atomic Properties 9
- 2. Slide 2 of 35 The Periodic Table and Some Atomic Properties Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9
- 3. Slide 3 of 35 9-1 Classifying the Elements: The Periodic Law and the Periodic Table 1869 Dimitri Mendeleev Lothar Meyer Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 3 of 35 When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically.
- 4. Slide 4 of 35 An illustration of the periodic law – variation of atomic volume with atomic number FIGURE 9-1 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 4 of 35
- 5. Slide 5 of 35 Mendeleev’s Periodic Table Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 5 of 35 — = 44 — = 72— = 68 — = 100
- 6. Slide 6 of 35 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 6 of 35
- 7. Slide 7 of 35 Atomic Number as the Basis for the Periodic Law Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 7 of 35 FIGURE 9-2 Schematic of an X-ray tube and Moseley’s X-ray spectra of several elements Moseley 1913 X-ray emission is explained in terms of transitions in which e- drop into orbits close to the atomic nucleus. Correlated frequencies to nuclear charges. ν = A (Z – b)2 Used to predict new elements (43, 61, 75) later discovered.
- 8. Slide 8 of 35 The Periodic table Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 8 of 35 Alkali Metals Alkaline Earths Transition Metals Halogens Noble Gases Lanthanides and Actinides Main Group Main Group
- 9. Slide 9 of 35 9-2 Metals and Nonmetals and Their Ions Metals Good conductors of heat and electricity. Malleable and ductile. Moderate to high melting points. Nonmetals Nonconductors of heat and electricity. Brittle solids. Some are gases at room temperature. Metalloids Metallic and non-metallic properties Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 9 of 35
- 10. Slide 10 of 35 Main-Group Metal Ions Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 10 of 35 Metals tend to lose electrons to attain noble gas electron configurations.
- 11. Slide 11 of 35 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 11 of 35
- 12. Slide 12 of 35 Main-Group Nonmetal Ions Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 12 of 35 Nonmetals tend to gain electrons to attain noble-gas electron configurations
- 13. Slide 13 of 35 9-3 Sizes of Atoms and Ions Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 13 of 35 FIGURE 9-3 Covalent, metallic, and ionic radii compared
- 14. Slide 14 of 35 The shielding effect and effective nuclear charge, Zeff FIGURE 9-6 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 14 of 35 Screen of electron charge from 10 core electrons (-) (-) (12+)
- 15. Slide 15 of 35 Screening and Penetration Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 15 of 35 FIGURE 9-4 Atomic radii
- 16. Slide 16 of 35 Ionic Radius Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 16 of 35 FIGURE 9-7 A comparison of atomic and ionic sizes
- 17. Slide 17 of 35 Ionic Radius Copyright © 2011 Pearson Canada Inc.Slide 17 of 35 General Chemistry: Chapter 9 Cations are smaller than the atoms from which they are formed. For isoelectronic cations, the more positive the ionic charge, the smaller the ionic radius. Anions are larger than the atoms from which they are formed. For isoelectronic anions, the more negative the charge, the larger the ionic radius.
- 18. Slide 18 of 35 Covalent and anionic radii compared FIGURE 9-8 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 18 of 35
- 19. Slide 19 of 35 A comparison of some atomic and ionic radii FIGURE 9-9 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 19 of 35
- 20. Slide 20 of 35 9-4 Ionization Energy Mg(g) → Mg+ (g) + e- I1 = 738 kJ Mg+ (g) → Mg2+ (g) + e- I2 = 1451 kJ Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 20 of 35 I = RH n2 Zeff 2 Ionization energies decrease as atomic radii increase.
- 21. Slide 21 of 35 First ionization energies as a function of atomic number FIGURE 9-10 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 21 of 35
- 22. Slide 22 of 35 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 22 of 35 I2 (Mg) vs. I3 (Mg) 7733 1451 I1 (Mg) vs. I1 (Al) 577.6 I1 (P) vs. I1 (S) 1012 999.6
- 23. Slide 23 of 35 9-5 Electron Affinity F(g) + e- → F- (g) EA = -328 kJ Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 23 of 35 F(1s2 2s2 2p5 ) + e- → F- (1s2 2s2 2p6 ) Li(g) + e- → Li- (g) EA = -59.6 kJ
- 24. Slide 24 of 35 Electron affinities of main-group elements FIGURE 9-11 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 24 of 35 Values are in kilojoules per mole for the process X(g) + e- X- (g).
- 25. Slide 25 of 35 Second Electron Affinities O(g) + e- → O- (g) EA = -141 kJ Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 25 of 35 O- (g) + e- → O2- (g) EA = +744 kJ
- 26. Slide 26 of 35 9-6 Magnetic Properties Diamagnetic atoms or ions: All e- are paired. Weakly repelled by a magnetic field. Paramagnetic atoms or ions: Unpaired e- . Attracted to an external magnetic field. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 26 of 35
- 27. Slide 27 of 35 Paramagnetism Copyright © 2011 Pearson Canada Inc.Slide 27 of 35 General Chemistry: Chapter 9
- 28. Slide 28 of 35 9-7 Periodic Properties of the Elements Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 28 of 35 FIGURE 9-12 Atomic properties and the periodic table – a summary
- 29. Slide 29 of 35 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 29 of 35 332266 ?? Three halogen elements FIGURE 9-13
- 30. Slide 30 of 35 Melting points of the third-period elements FIGURE 9-14 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 30 of 35
- 31. Slide 31 of 35 Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 31 of 35
- 32. Slide 32 of 35 Reducing Abilities of Group 1 and 2 Metals Copyright © 2011 Pearson Canada Inc.Slide 32 of 35 General Chemistry: Chapter 9 2 K(s) + 2 H2O(l) → 2 K+ + 2 OH- + H2(g) Ca(s) + 2 H2O(l) → Ca2+ + 2 OH- + H2(g) I1 = 419 kJ I1 = 590 kJ I2 = 1145 kJ
- 33. Slide 33 of 35 Oxidizing Abilities of the Halogen Elements (Group 17) Copyright © 2011 Pearson Canada Inc.Slide 33 of 35 General Chemistry: Chapter 9 2 Na + Cl2 → 2 NaCl Cl2 + 2 I- → 2 Cl- + I2
- 34. Slide 34 of 35 Acid-Base Nature of Element Oxides Copyright © 2011 Pearson Canada Inc.Slide 34 of 35 General Chemistry: Chapter 9 Basic oxides or base anhydrides: Li2O(s) + H2O(l) → 2 Li+ (aq) + 2 OH- (aq) Acidic oxides or acid anhydrides: SO2 (g) + H2O(l) → H2SO3(aq) Na2O and MgO yield basic solutions Cl2O, SO2 and P4O10 yield acidic solutions SiO2 dissolves in strong base, acidic oxide.
- 35. Slide 35 of 35 End of Chapter Questions Think of a problem like a root system: Each branching is a decision. The answer is at the tip of one of the rootlets. Copyright © 2011 Pearson Canada Inc.General Chemistry: Chapter 9Slide 35 of 35 Answer a b c d fe hg
Editor's Notes
- In this chapter, you will learn how to deduce and write chemical formulas and how to use the information incorporated into chemical formulas. The chapter ends with an overview of the relationship between names and formulas—chemical nomenclature.
- A scanning tunneling microscope image of 48 iron atoms adsorbed onto a surface of copper atoms. The iron atoms were moved into position with the tip of the scanning tunneling microscope in order to create a barrier that forced some electrons of the copper atoms into a quantum state seen here as circular rings of electron density. The colors are from the computer rendering of the image. In this chapter we discuss the periodic table and the properties of atoms and ions. In this chapter, we will use the table as a backdrop for a discussion of some properties of elements, including atomic radii, ionization energies, and electron affinities. These atomic properties also arise in the discussion of chemical bonding in the following two chapters, and the periodic table itself will be our indispensable guide throughout much of the remainder of the text.
- In 1869, Dmitri Mendeleev and Lothar Meyer independently proposed the periodic law: When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically. Meyer based his periodic law on the property called atomic volume—the atomic mass of an element divided by the density of its solid form. We now just call this property molar volume.
- This adaptation of Meyer’s 1870 graph plots atomic volumes against atomic numbers. Of course, a number of elements, such as the noble gases, were undiscovered in Meyer’s time. The graph shows peaks at the alkali metals (Li, Na, K, and so on). Nonmetals fall on the ascending portions of the curve and metals at the peaks, on the descending portions, and in the valleys. Meyer presented his results as a graph of atomic volume against atomic mass. Now it is customary to plot his results as molar volume against atomic number, as shown. Notice how high atomic volumes recur periodically for the alkali metals Li, Na, K, Rb, and Cs. Later, Meyer examined other physical properties of the elements and their compounds, such as hardness, compressibility, and boiling points, and found that these also vary periodically.
- Mendeleev’s work attracted more attention than Meyer’s for two reasons: He left blank spaces in his table for undiscovered elements, and he corrected some atomic mass values. The blanks in his table came at atomic masses 44, 68, 72, and 100 for the elements we now know as scandium, gallium, germanium, and technetium. Two of the atomic mass values he corrected were those of indium and uranium. Meyer’s work was seen to support Mendeleev’s and convinced the scientific community to accept this approach.
- Other additions to the table included the Noble gases discoved by William Ramsey. Mendeleev placed certain elements out of order-he assumed that errors had been made in the atomic masses, but it is clear that some elements remain out of order. Moseley changed that with x-ray spectra.
- 43, 61 and 75 were discovered in 1937, 1945 and 1925 respectively. Also proved the periodic law in the region from Z = 13 to 79, and that there could be NO other elements in this region. (a) A heated filament emits electrons by a process called thermionic emission. The electrons are accelerated by a high voltage, and collide with the metal target. The highly energetic electrons ionize electrons from the inner shells of the metal atoms of the target. Subsequently, electrons from higher orbitals drop down to occupy the vacancies and in doing so, emit X-ray photons that correspond to the energy difference between the two orbitals. (b) In this photograph from Moseley’s 1913 paper, you can see two lines for each element, beginning with Ca at the top. With each successive element, the lines are displaced to the left, the direction of increasing X-ray frequency in these experiments. Where more than two lines appear, the sample contained one or more other elements, or impurities. Notice, for example, that one line in the Co spectrum matches a line in the Fe spectrum, and another matches a line in the Ni spectrum. Brass, which is an alloy of copper and zinc, shows two lines for Cu and two for Zn.
- In the periodic table the vertical groups bring together elements with similar properties. The horizontal periods of the table are arranged in order of increasing atomic number from left to right. The groups are numbered at the top, and the periods at the extreme left. The first two groups—the s block—and the last six groups—the p block—together constitute the main-group elements. Because they come between the s block and the p block, the d block elements are known as the transition elements. The f block elements, sometimes called the inner transition elements, would extend the table to a width of 32 members if incorporated in the main body of the table. The table would generally be too wide to fit on a printed page, and so the f block elements are extracted from the table and placed at the bottom. The 15 elements following barium Z=56 are called the lanthanides, and the 15 following radon Z=88 are called the actinides.
- Discuss Transition Metals on This slide when transition metal atoms ionize, the ns subshell is emptied. Ti is [Ar]3d24s2,but Ti2+ is [Ar]3d2 and Ti4+ is [Ar]. The 3d subshell in Fe3+ is half-filled, a fact that helps to account for the observed ease of oxidation of iron(II) to iron(III) compounds. Electron configurations with half-filled or filled d or f subshells have a special stability, and a number of transition metal ions have such configurations.
- Two valence electrons (blue) are attracted to the nucleus of a Mg atom. The atom’s nuclear charge is screened by the 10 core electrons (gray), but not perfectly. The valence electrons also screen each other somewhat. The result is an effective nuclear charge, Zeff, closer to 3+ than to 2+.
- Radii increase down a group. Radii decrease across a period in the main group (Zeff increases across main group elements). Radii in Transition metals remain fairly constant except for a few spikes. Electrons go into an inner shell, thus participate in shielding the outer shell electrons from the increasing Zeff.
- metallic and ionic radii are shown
- Knowdledge of atomic and ionic radii can be used to varycertain physical porperties. Ex Na+ and Ca2+ ions. Glass is brittle and breaks easily. Replace surface Na+ by K+ and glass becomes shatter resistant. Cr3+ in Al2O3 (about 1%) gives beautiful red colour (Ruby).
- In the case of magnesium, for example, in the second ionization, the electron, once freed, has to move away from an ion with a charge of 2+ (Mg2+). More energy must be invested than for a freed electron to move away from an ion with a charge of 1+ (Mg+) . This is a direct consequence of Coulomb’s law, which states, in part, that the force of attraction between oppositely charged particles is directly proportional to the magnitudes of the charges. Across a period, as Zeff increases and the valence-shell principal quantum number n remains constant, the ionization energy should increase. And down a group, as n increases and Zeff increases only slightly, the ionization energy should decrease. Thus, atoms lose electrons more easily (become more metallic) as we move from top to bottom in a group of the periodic
- Because their electron configurations are so stable, more energy is required to ionize noble gas atoms than to ionize atoms of the elements immediately preceding or following them. The maxima on the graph come at the atomic numbers of the noble gases. The alkali metals are the most easily ionized of all groups. The minima in the graph come at their atomic numbers.
- Removing the third electon from Mg causes a large jump in I. I1 of Al less than Mg because s- electron is removed from Mg and p-electron is removed from Al. I1 of S is less than that of P. This is due to e--e- repulsion of the fourth electron.
- Notice that the process above is exothermic, meaning that energy is given off when an F atom gains an electron. Electron affinity, EA, can be defined as the enthalpy change ΔHea , that occurs when an atom in the gas phase gains an electron. According to this definition, the electron affinity of fluorine is a negative quantity.
- It is more difficult to make generalizations about electron affinities than about ionization energies. The smaller atoms to the right of the periodic table (for example, group 17) tend to have large, negative electron affinities.* Electron affinities tend to become less negative in progressing toward the bottom of a group, with the notable exception of the second-period members of groups 15, 16, and 17 (namely, N, O, and F). It is likely that for these small atoms, an incoming electron encounters strong repulsive forces from other electrons in the atom and is thereby not as tightly bound as we might otherwise expect.
- In considering the gain of a second electron by a nonmetal atom, we encounter positive electron affinities. Here the electron to be added is approaching not a neutral atom, but a negative ion. There is a strong repulsive force between the electron and the ion, and the energy of the system increases. Thus, for an element like oxygen, the first electron affinity is negative and the second is positive. The high positive value of EA2 makes the formation of gaseous O2- seem very unlikely. The O2- ion can exist, however, in ionic compounds, such as MgO(s), where formation of the ion is accompanied by other energetically favorable processes.
- Manganese has a paramagnetism corresponding to five unpaired electrons, which is consistent with the electron configuration. When a manganese atom loses two electrons, it becomes the Mn2+ ion which is paramagnetic, and the strength of its paramagnetism corresponds to five unpaired electrons. When a third electron is lost to produce Mn3+, the ion has a paramagnetism corresponding to four unpaired electrons. The third electron lost is one of the unpaired 3d electrons.
- Metallic character corresponds to conductance of heat and electricity.
- Average is 349 for mp Average is
- Melting involves destruction of the orderly arrangement of atoms or molecules in a crystalline solid. Melting point temperature depends on the strength of the attractive forces between atoms of molecules in the solid. Na, Mg, Al Metallic bonds, Si Covalent bonds, strong interatomic forces. P4, S8, Cl2 Discrete molecules, forces become weaker as you go across.
- K is representative of reactivity of Group 1, expect that a lower ionization energy should react more vigerously. Ca is representative of reactivity of Group 2, expect that a higher ionization energy should react more slowly. Ionization energy alone is an oversimlification. If differences in I are small then other factors must be taken into account.
- Expect compounds with high electron affinity to be good oxidizing agents. So halogens should react vigorously with Na. Cl has higher electron affinity than I, therefore the reaction should lie to the right.