1 / 37

Chapter 9. Periodicity: The Periodic Table and Atomic Properties

Chapter 9. Periodicity: The Periodic Table and Atomic Properties. Periodic Law: when arranged by atomic mass, the elements exhibit a periodic recurrence of similar properties. Chemical reactivity is governed by the electronic structure of a species (ie. an atom).

jeanette-lara
Download Presentation

Chapter 9. Periodicity: The Periodic Table and Atomic Properties

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 9. Periodicity: The Periodic Table and Atomic Properties Periodic Law: when arranged by atomic mass, the elements exhibit a periodic recurrence of similar properties Chemical reactivity is governed by the electronic structure of a species (ie. an atom). • in this chapter we will learn how the organization of the periodic table is explained by the new quantum mechanical model of the atom in other words, • how does the electron configuration of an element, the distribution of electrons within the atom’s orbitals, relate to the atom’s physical and chemical properties

  2. 1871 1886 Property of Germanium (Ge, E) Predicted by Mendeleev (E) Actual (Ge)

  3. Trends in Some Key Periodic Atomic Properties • all physical and chemical properties of the elements are based ultimately on the electron configuration of their atoms Trends in Atomic Size • we define the size of atoms by how closely one atom lies next to another • the metallic radius is one half the distance between the nuclei of adjacent atoms in a crystal of the element • the covalent radius is one half the distance between the nuclei of identically bonded atoms 186 pm 157 pm Na(s) Na2(g)

  4. Trends in Atomic Size Main Group • moving down a group each member of the new period has another level of electrons (n increases). The larger n, the farther from the nucleus the outer electrons spend time and the radius is larger. • moving across a period, electrons are added to the outer level. However, the nuclear charge also increases. The effective nuclear charge felt by the outer electrons increases and, so the atomic size decreases….

  5. Remember Screening… The effective nuclear charge, Zeff, is Z = nuclear charge S = screening factor Ns = # electrons in the same shell N1 = # electrons in n-1 shell No = # other electrons (n-2 shell or lower) ie. for P 1s22s22p63s23p3 ie. for S 1s22s22p63s23p4

  6. Transition Elements Group Co: 1s22s22p63s23p63d74s2 • the trends are good for the first few across a period but the size of the atoms remain constant after that because of shielding by the added inner d electrons counteracting the usual increase in Zeff. Ni: 1s22s22p63s23p63d84s2 2.20 2.85 3.00 3.15 3.30 2.95 3.60 3.75 3.90 4.05 3.70 4.35 5.0 5.65 6.30

  7. the np electrons are further from the nucleus than the (n-1)d electrons and therefore are shielded from the nucleus by all the electrons in the d orbitals, but not fully (only by 0.85). Therefore there is a much larger size decrease going from Ca to Ga than from Mg to Al in the main group. Mg: 1s22s22p63s2 Al: 1s22s22p63s23p1 Ca: 1s22s22p63s23p64s2 Ga: 1s22s22p63s23p63d104s24p1

  8. Trends in Ionization Energy (Ionization Potential, IP(eV) • the energy required to remove an electron from an atom, molecule or ion • or the energy released an electron is attached to a cation the first ionization energy, (IE1) the 2nd ionization energy, (IE2)

  9. eg. Name the period 3 elements with the following ionization energies (in kJ mol-1) and write their electron configurations. • IE1 IE2 IE3 IE4 IE5 IE6 • 1012 1903 2910 4956 6278 22230 • 577 1816 2744 11576 14829 18375 Solution. First three ionization energies of Be.

  10. Which of numbers in the periodic table below corresponds to the group of elements in which the 4d orbitals are being filled? 1 2 3 4 5

  11. Which of the following is the condensed electron configuration for chlorine? 1. [Ne]3p7 2. [Ne]3s23p5 Cl2(g) 3. [Ne]3s23p6 4. [Ne]3s23d5 Cl2(l) 5. [Ne]3s23p33d2

  12. In general the ionization energy increases going across a period. However, as seen to the right, the ionization energy of Al is considerably smaller than that of Mg. The best explanation for this is 1. Al is a smaller atom than Mg. 2. s electrons penetrate closer to the nucleus and therefore shield p electrons more effectively. • the electrons in Mg are farther from the nucleus than in Al. 4. Al is a larger atom than Mg since its electrons are less tightly held to the nucleus.

  13. Your text is uses the engineers notation for the sign of EA, as does most texts Trends in Electron Affinity • the energy required to remove an electron from a singly charged negative ion • or the energy released when an electron is attached to an atom 0.0 -18 -29 -21 -35 • the trends in Ea are not as well defined as they are for size and IE, likely due to electron-electron repulsion -186 -39 -146 -41 -46 -41

  14. all EAn’s for negative ions are negative due to repulsion between the ion and electron • from this, it seems that O2- is not likely to be formed, but in ionic compounds like MgO, O exists as O2-. In fact gaseous O2- will not be formed by adding an electron. • O2- or other multiply negatively charged ions are formed in the presence of other energetically favorable processes (hang on till Chapter 13).

  15. O+ Kr Kr+ O Below is an example of a charge exchange reaction between atomic oxygen and the krypton cation. These reactions are possible if they are exothermic. The ionization energies of Kr and O are 1350 and 1312 kJ mol-1, respectively. Is the charge exchange reaction shown below exothermic or endothermic? + +

  16. Trends in Metallic Behavior Metalloids Non-metals, Metals • typically not shiny • relatively low melting points • poor thermal and electrical conductors • tend to gain electrons in a chemical reaction • shiny • high melting points • good conductors of heat and electricity • lose electrons in a chemical reaction

  17. The change in metallic behavior across period 3 and down group 15. N and P are non-metals forming 3- ions in ionic compounds IE1 As and Sb are metalloids and don’t actually form ions readily Bi is a metal forming 3+ ions in its compounds

  18. Reducing Abilities of Group 1 and 2 Metals Group 1 Group 2 Li Be Na Mg K Ca Rb

  19. Oxidizing Abilities of the Halogens (group 17) general oxidizing ability of elements (increase as metallic character decreases)

  20. Acid-Base Behavior of the Element Oxides blue: act as bases red: act as acids

  21. Which of the following reactions would you expect to be the most vigorous?

  22. The Non-Reactivity of the Group 18 Elements (Rare Gases) • the rare gases have a completed “octet” and for many years after their discovery no compounds containing rare gases were known • in 1962 Bartlett and Lohmann prepared a yellow crystalline solid that they thought was XePtF6 (it was actually Xe(PtF6)n where 1<n<2) • XeF2 • XeF4 • XeOF2 • XeF6 • XeO3 • XeO4 • H4XeO6 • KrF2

  23. many other unstable or reactive rare-gas containing ions and neutrals have been prepared under “matrix isolation” conditions.

  24. catalysis by rare gases!

  25. Electron Configurations of Ions Main Group Ions • main group elements gain or lose electrons to attain a rare gas electron configuration For example, the electron configurations of O and its common ion are, O: [He]2s22p4 O2-: [He]2s22p6 = [Ne] and for Mg and Mg2+ Mg: [Ne]2s2 Mg2+: [Ne]

  26. the larger metal ions in period 4 and 5 of groups 13, 14 and 15 would have to lose far too many electrons to attain a rare-gas electron configuration but do form positive ions. for example, Sn: [Kr]5s24d105p2 can lose 4 electrons Sn4+: [Kr]4d10 a pseudo-rare gas configuration of electrons more commonly, Sn loses just the p electrons Sn2+: [Kr]5s24d10

  27. for some elements, such as carbon, it is too energetically costly to attain a rare-gas configuration by gaining or losing 4 electrons. These elements, we will see, attain a rare-gas electron configuration by “sharing” electrons when they form compounds.

  28. Ionic vs Atomic Size

  29. Which of the following has the smallest radius? 1. Li+ 2. Li 3. O2- 4. Be2+ 5. Ne

  30. 3d 3d 3d 3d 3d 3d 3d 3d [Ar] [Ar] 4s 4s [Ar] [Ar] 4s 4s [Ar] 4s [Ar] [Ar] 4s 4s [Ar] 4s Transition Metal Ions • transition metal ions rarely attain a rare gas configuration since the energy costs would be too high to lose many electrons • in Period 4 Sc forms Sc3+ and Ti forms Ti4+ in some compounds • typically transition metals form more than one ion, by losing all of their ns electrons and some of their (n-1)d electrons • they lose their ns electrons first…first in, first out for transition metal ions.(why many metals form 2+ ions) Mn: [Ar]3d54s2 Mn2+: [Ar]3d5 Fe2+: [Ar]3d6 Fe: [Ar]3d64s2 Fe3+: [Ar]3d5 Cu: [Ar]3d104s1 Cu+: [Ar]3d10 Cu2+: [Ar]3d9

  31. How do we know the electron configuration of atoms and ions?? One piece of evidence can come from the magnetic properites: Diamagnetic and Paramagnetic • a species which has unpaired electrons is paramagnetic • a species which has only paired electrons is diamagnetic The extent to which a species is attracted into the magnetic field is proportional to the number of unpaired electrons

  32. [Ar] Ti: [Ar]3d24s2 Ti2+: [Ar]3d2 3d 4s 3d 4s 3d 3d Fe: [Ar]3d64s2 Fe3+: [Ar]3d5 [Ar] 4s [Ar] 4s Let’s look at Ti, Both Ti and compounds of Ti2+ are paramagnetic, to the same extent supporting the conclusion that these elements lose their s electrons first When iron forms Fe3+, there is an increase in paramagnetism because it loses its 4s electrons and a 3d electron. If it lost only 3d electrons, the paramagnetism would decrease.

  33. eg. Use condensed electron configurations to write the reaction for the formation of each transition metal ion, and predict whether the ion is paramagnetic. a) Mn2+ b) Cr3+ c) Hg2+ d) Co3+

More Related