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Exam R Nov 18 10-11:30 am

Exam R Nov 18 10-11:30 am. Review session W Nov 17 11-noon CF 316. Specify if each pair has chemical properties that are similar (  ) or not similar (X) : 1. Cl and Br 2. P and S 3. O and S. Concept Check.

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Exam R Nov 18 10-11:30 am

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  1. Exam R Nov 1810-11:30 am Review session W Nov 17 11-noon CF 316

  2. Specify if each pair has chemical properties that are similar () or not similar (X): 1. Cl and Br 2. P and S 3. O and S Concept Check

  3. Atomic radius is plotted against atomic number in the graph below. Note the regular (periodic) variation.

  4. A representation of atomic radii is shown below.

  5. 34 Se 35 Br 52 Te • Refer to a periodic table and arrange the following elements in order of increasing atomic radius: Br, Se, Te. Te is larger than Se. Se is larger than Br. Br < Se < Te

  6. Atomic radii of main-group elements

  7. Understanding “periodicity” Trends within a period what is changing? what is constant? Trends within a group what is changing? what is constant? 11 Na 12 Mg 11 Na 19 K

  8. Screening effect of inner electrons Consider Mg [Ne]3s2 Core electrons of [Ne] “shield” valence electrons

  9. Periodicity – trends within a given period Trends within a period what is changing? what is constant? 11 Na 12 Mg 13 Al

  10. Effective nuclear charge (Z*) Z* = nuclear charge experienced by the outermost electrons. Estimate Z* by  [ Z - (# of shielding electrons)] (Z = total number of electrons) Z* increases across a period

  11. First Ionization Energy (first ionization potential) • The minimum energy needed to remove the highest-energy (outermost) electron from a neutral atom in the gaseous state, thereby forming a positive ion

  12. These trends and reversals are visible in the graph of ionization energy versus atomic number.

  13. The size of each sphere indicates the size of the ionization energy in the figure below.

  14. Electrons can be successively removed from an atom. Each successive ionization energy increases, because the electron is removed from a positive ion of increasing charge. • A dramatic increase occurs when the first electron from the noble-gas core is removed.

  15. Left of the line, valence shell electrons are being removed. Right of the line, noble-gas core electrons are being removed.

  16. Trends • Going down a group, first ionization energy decreases. • This trend is explained by understanding that the smaller an atom, the harder it is to remove an electron, so the larger the ionization energy.

  17. Generally, ionization energy increases with atomic number. • Ionization energy is proportional to the effective nuclear charge divided by the average distance between the electron and the nucleus. Because the distance between the electron and the nucleus is inversely proportional to the effective nuclear charge, ionization energy is inversely proportional to the square of the effective nuclear charge.

  18. 33 As 35 Br 51 Sb • Refer to a periodic table and arrange the following elements in order of increasing ionization energy: As, Br, Sb. Sb is larger than As. As is larger than Br. Ionization energies: Sb < As < Br

  19. Small deviations occur between Groups IIA and IIIA and between Groups VA and VIA. • Examining the valence configurations for these groups helps us to understand these deviations: • IIA ns2 • IIIA ns2np1 • VA ns2np3 • VIA ns2np4 Ittakes less energy to remove the np1 electron than the ns2 electron. It takes less energy to remove the np4 electron than the np3 electron.

  20. Electrons can be successively removed from an atom. Each successive ionization energy increases, because the electron is removed from a positive ion of increasing charge. • A dramatic increase occurs when the first electron from the noble-gas core is removed.

  21. Left of the line, valence shell electrons are being removed. Right of the line, noble-gas core electrons are being removed.

  22. 33 As 35 Br 51 Sb • Refer to a periodic table and arrange the following elements in order of increasing ionization energy: As, Br, Sb. Sb is larger than As. As is larger than Br. Ionization energies: Sb < As < Br

  23. Electron affinity (E.A.) • The energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion • A negative energy change (exothermic) indicates a stable anion is formed. The larger the negative number, the more stable the anion. Small negative energies indicate a less stable anion. • A positive energy change (endothermic) indicates the anion is unstable.

  24. The electron affinity is > 0, so the element must be in Group IIA or VIIIA. The dramatic difference in ionization energies is at the third ionization. The element is in Group IIA.

  25. Broadly speaking, the trend is toward more negative electron affinities going from left to right in a period.

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