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Periodic Relationships Among the Elements

Periodic Relationships Among the Elements. Chapter 8. Periodic Table - In the Beginning. 1. A necessary prerequisite to the construction of the periodic table was the discovery of the individual elements.

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Periodic Relationships Among the Elements

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  1. Periodic Relationships Among the Elements Chapter 8

  2. Periodic Table - In the Beginning 1 • A necessary prerequisite to the construction of the periodic table was the discovery of the individual elements. • Although elements such as gold, silver, tin, copper, lead, and mercury have been known since antiquity, the first scientific discovery of an element occurred in 1649 when Hennig Brand discovered phosphorous. • By 1869, a total of 63 elements had been discovered. As the number of known elements grew, scientists began to recognize patterns in properties and began to develop classification schemes.

  3. Periodic Table – Law of Triads 1 • In 1817 Johann Dobereiner noticed that the atomic weight of strontium fell midway between the weights of calcium and barium, elements possessing similar chemical properties. • In 1829, after discovering the halogen triad composed of chlorine, bromine, and iodine and the alkali metal triad of lithium, sodium and potassium he proposed that nature contained triads of elements the middle element had properties that were an average of the other two members when ordered by the atomic weight (the Law of Triads).

  4. Periodic Table – Law of Triads 1 • This new idea of triads became a popular area of study. Between 1829 and 1858 a number of scientists (Jean Baptiste Dumas, Leopold Gmelin, Ernst Lenssen, Max von Pettenkofer, and J.P. Cooke) found that these types of chemical relationships extended beyond the triad. • During this time fluorine was added to the halogen group; oxygen, sulfur, selenium and tellurium were grouped into a family while nitrogen, phosphorus, arsenic, antimony, and bismuth were classified as another. • Unfortunately, research in this area was hampered by the fact that accurate values of were not always available.

  5. First Attempts At Designing a Periodic Table 1 • in 1862 French geologist, A.E.Beguyer de Chancourtois transcribed a list of the elements positioned on a cylinder in terms of increasing atomic weight. • When the cylinder was constructed so that 16 mass units could be written on the cylinder per turn, closely related elements were lined up vertically. • De Chancourtois was first to recognize that elemental properties reoccur every seven elements, and using this chart, he was able to predict the stoichiometry of several metallic oxides. Unfortunately, his chart included some ions and compounds in addition to elements.

  6. Law of Octaves 1 • John Newlands, an English chemist, wrote a paper in 1863 which classified the 56 established elements into 11 groups based on similar physical properties, noting that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. • In 1864 Newlands published his version of the periodic table and proposed the Law of Octaves (by analogy with the seven intervals of the musical scale). This law stated that any given element will exhibit analogous behavior to the eighth element following it in the table.

  7. Father of Periodic Table 1 • There has been some disagreement about who deserves credit for being the "father" of the periodic table, the German Lothar Meyer or the Russian Dmitri Mendeleev. • Both chemists produced remarkably similar results at the same time working independently of one another.

  8. Father of Periodic Table 1 • Meyer's 1864 textbook included a rather abbreviated version of a periodic table used to classify the elements. This consisted of about half of the known elements listed in order of their atomic weight and demonstrated periodic valence changes as a function of atomic weight. • In 1868, Meyer constructed an extended table which he gave to a colleague for evaluation. • Unfortunately for Meyer, Mendeleev's table became available to the scientific community via publication (1869) before Meyer's appeared (1870).

  9. Periodic Table – Noble Gases 1 • In 1895 Lord Rayleigh reported the discovery of a new gaseous element named argon which proved to be chemically inert. This element did not fit any of the known periodic groups. • In 1898, William Ramsey suggested that argon be placed into the periodic table between chlorine and potassium in a family with helium, despite the fact that argon's atomic weight was greater than that of potassium. This group was termed the "zero" group due to the zero valiancy of the elements. Ramsey accurately predicted the future discovery and properties neon.

  10. The Modern Periodic Table 1 • The last major changes to the periodic table resulted from Glenn Seaborg's work in the middle of the 20th Century. Starting with his discovery of plutonium in 1940, he discovered all the transuranic elements from 94 to 102 using the particle accelerator at University of California, Berkeley. • He reconfigured the periodic table by placing the actinide series below the lanthanide series. • In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. • Element 106 has been named seaborgium (Sg) in his honor.

  11. The Periodic Law and the Periodic Table 1 • Periodic Law - the physical and chemical properties of the elements are periodic functions of their atomic numbers.

  12. When the Elements Were Discovered

  13. ns2np6 ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 4f 5f Ground State Electron Configurations of the Elements

  14. Classification of the Elements

  15. Electron Arrangement and the Periodic Table 3 • Electron configuration - describes the arrangement of electrons in atoms. • The electron arrangement is the primary factor in understanding how atoms join together to form compounds. • Valance electrons - the outermost electrons. • These are the electrons involved in chemical bonding.

  16. Valance Electrons • For the representative elements: • The number of valance electrons is the group number. • The period number gives the energy level (n) of the valance shell. • For an atom of fluorine, how many valance electrons does it have and what is the energy level of these electrons? • Fluorine has 7 electrons in the n=2 level

  17. Let’s look at fluorine more closely. • What is the total number of electrons in fluorine? • The atomic number is 9. It therefore has 9 protons and 9 electrons. • If there are 7 electrons in the valance shell, (with n = 2 energy level) where are the other two electrons? • In the n = 1 energy level. This level holds two and only two electrons.

  18. Isoelectronic - they have the same electron configuration (same number of electrons) • Nonmetallic elements tend to form negatively charged ions called anions. • Nonmetals tend to gain electrons so they become isoelectronic with its nearest noble gas neighbor. O O2- + 2e- [He]2s22p4 [He]2s22p6 or [Ne]

  19. Electron Configurations of Cations and Anions Of Representative Elements Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne]

  20. -1 -2 -3 +1 +2 +3 Cations and Anions Of Representative Elements

  21. What neutral atom is isoelectronic with H- ? Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne H-: 1s2 same electron configuration as He

  22. Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5

  23. Trends in the Periodic Table • We will look at the following trends • In effective nuclear charge • in atomic size • in ion size • in ionization energy • in electron affinity

  24. Periodic Properties • Two factors determine the size of an atom. • One factor is the principal quantum number, n. The larger is “n”, the larger the size of the orbital. • The other factor is the effective nuclear charge, which is the positive charge an electron experiences from the nucleus minus any “shielding effects” from intervening electrons.

  25. Core Radius (pm) Z Zeff Na 11 10 1 186 12 10 2 160 Mg Al 13 10 3 143 Si 14 10 4 132 Effective nuclear charge (Zeff) is the “positive charge” felt by an electron. Zeff = Z - s 0 < s < Z (s = shielding constant) Zeff Z – number of inner or core electrons

  26. Effective Nuclear Charge (Zeff) increasing Zeff increasing Zeff

  27. Atomic Size 7 • 1. The size of the atoms increases from top to bottom. • This is due to the valance shell being higher in energy and farther from the nucleus. • 2. The size of the atoms decreases from left to right. • This is due to the increase in magnitude of positive charge in nucleus. The nuclear charge pulls the electrons closer to the nucleus.

  28. Atomic Radii

  29. Variation in Size of Atoms

  30. Comparison of Atomic Radii with Ionic Radii

  31. Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.

  32. Ion Size 7 • Cations are always smaller than their parent atom. • This is due to more protons than electrons. The extra protons pulls the remaining electrons closer. • Which would be smaller, Fe2+ or Fe3+? • Fe3+ • This size trend is also due to the fact that it is the outer shell that is lost.

  33. Anions are always larger than their parent atom. • This is due to the fact that anions have more electrons than protons.

  34. The Radii (in pm) of Ions of Familiar Elements

  35. Ionization Energy 7 8 • Ionization energy - The energy required to remove an electron from an isolated atom. ionization energy + Na  Na+ + e- • The magnitude of ionization energy correlates with the strength of the attractive force between the nucleus and the outermost electron. • The lower the ionization energy, the easier to form a cation.

  36. I1 + X (g) X+(g) + e- I2 + X+(g) X2+(g) + e- I3 + X2+(g) X3+(g) + e- Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. I1 first ionization energy I2 second ionization energy I3 third ionization energy I1 < I2 < I3

  37. Variation of the First Ionization Energy with Atomic Number Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell

  38. Ionization decreases down a family because the outermost electrons are farther from the nucleus. • Ionization increases across a period because the outermost electrons are more tightly held. • Why do you think that the noble gases would be so unreactive?

  39. Increasing First Ionization Energy Increasing First Ionization Energy General Trend in First Ionization Energies

  40. Electron Affinity 7 8 • Electron Affinity - The energy change when a single electron is added to an isolated atom. Br + e- Br- + energy • Electron affinity gives information about the ease of anion formation. • Large electron affinity indicates an atom becomes more stable as it forms an anion.

  41. X (g) + e- X-(g) F (g) + e- X-(g) O (g) + e- O-(g) Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. DH = -328 kJ/mol EA = +328 kJ/mol DH = -141 kJ/mol EA = +141 kJ/mol

  42. Electron Affinity = -349 kJ/mol Periodic Properties • The electron affinity is the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion. • Electron Affinity • For a chlorine atom, the first electron affinity is illustrated by:

  43. Periodic Properties • Electron Affinity • The more negative the electron affinity, the more stable the negative ion that is formed. • Broadly speaking, the general trend goes from lower left to upper right as electron affinities become more negative.

  44. Variation of Electron Affinity With Atomic Number (H – Ba)

  45. E.A. generally decreases down a group. • E.A. generally increases across a period.

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