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Chemistry 101 : Chap. 7

Chemistry 101 : Chap. 7. Periodic Properties of the Elements. Development of the Periodic Table (2) Effective Nuclear Charge (3) Sizes of Atoms and Ions (4) Ionization Energy (5) Electron Affinity Reading Assignment : 7.6-7.8. Periodic Table.

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Chemistry 101 : Chap. 7

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  1. Chemistry 101 : Chap. 7 Periodic Properties of the Elements • Development of the Periodic Table • (2) Effective Nuclear Charge • (3) Sizes of Atoms and Ions • (4) Ionization Energy • (5) Electron Affinity • Reading Assignment : 7.6-7.8

  2. Periodic Table • Mendeleev ordered known elements according to their weight D. Mendelejeff, Zeitscrift für Chemie12, 405-406 (1869)  In the modern version, elements are ordered according to atomic number Dimitri Mendeleev (1834-1907)  Elements within a same vertical group have similar properties

  3. Periodic Table Halogen Alkaline earth metal Transition metals Alkali metal rare gas

  4. Periodic Properties  Properties (e.g. reactivity) of atoms depends on…  Number of electrons  Electron configuration  How tightly electrons are bound to nucleus

  5. Periodic Table and Electron Configurations

  6. + + + + + + + - - - - - - - Effective Nuclear Charge  Effective nuclear charge (Zeff) is used to measure how tightly (outer-shell) electrons are bound to nucleus Zeff = # of proton - # of core electrons Valence electrons see nucleus and core electrons as a single unit core  Valence electrons experience approximately the total charge of protons and core electrons valence electrons

  7. Effective Nuclear Charge  Example : What is the effective charge of 2p electron of N ?  Example : What is the effective charge of 4s electron of Ca?

  8. Effective Nuclear Charge Effective charge, Zeff,, increases with the atomic numbers within a period. 11Na 12Mg 13Al 14Si …. +11-10=+1 +12-10=+2 +13-10=+3 +14-10=+4 Zeff increases along a period  # of protons increases, but the # of core electrons stay the same along period. Valence electrons are more strongly bound as atomic number increases within a period

  9. Effective Nuclear Charge  Example : Which electron is more tightly bound: 2p electron of Ne or 3s electron of Na? Which electron would be easier to remove?

  10. Atomic Radius Atomic radii decrease with the atomic numbers within a period. Increasing Zeff along a period attract valence electrons more strongly, making the atom more compact Fluorine Lithium

  11. Atomic Radius Atomic radii increase with the atomic numbers within a group. +11 – 10 = +1 11Na +19 – 18 = +1  Zeff essentially remains constant 19K +37 – 36 = +1 37Rb However, more main shells are added and principal quantum number of valence electrons increase

  12. Atomic Radius atomic radius increases atomic radius increases atomic radius increases

  13. Ion Radius cations are always smaller than their parent atoms anions are always larger than their parent atoms

  14. Ion Radius  Example : Which of the following ions has the largest radius? (1) S, S2-, O2- (2) O2-, Na+, Al3+

  15. Ionization Energy  Ionization Energies (IE) : Energy required to remove electrons from an atom in the gas phase First Ionization Energy : Energy required to remove one electron from a neutral atom in the gas phase IE1 = 495 kJ/mol Na+ (g) + e- Na (g) Second Ionization Energy : Energy required to remove second electron from an atom in the gas phase IE2 = 4562 kJ/mol Na2+ (g) + e- Na+ (g)

  16. Ionization Energy First IE increases in general with atomic numbers within a period First IE decreases with atomic numbers within a group Increasing IE

  17. Ionization Energy  First IE increases with atomic numbers in a period because .. Effective nuclear charge increases  decreasing the distance from the nucleus  stronger interaction between valence electron and nucleus  First IE decreasewith atomic numbers in a group because .. Atomic radius increases with little change in effective nuclear charge  weaker interaction between valence electron and nucleus

  18. Ionization Energy • Example : Why there is a huge gap between the 5th and 6th ionization energy of nitrogen? Ionization E (kJ/mol) Ionization number

  19. Ionization Energy  Example : Why there are irregularities in the first IE within a period? 2p4 2s2 2p1

  20. Electron Affinity • Electron Affinity (Ea) : Energy change that occurs when an electron is added to a neutral atom in the gas phase DEa = - 349 kJ/mol Cl (g) + e- Cl- (g) 100 pm [Ne]3s23p5 167 pm [Ne]3s23p6 Energy is released ! More negative Ea means greater attraction between a given atom and an added electron

  21. Electron Affinity Added electron goes to new shell Added electron goes to new subshell Added electron leads to noble gas configuration

  22. Supplementary Material • Metals : Tend to loose electrons Small ionization energy (IE) Reactivity of metal is related to the IE  smaller IE = more reactive Which one is the most reactive: Na, Mg or K? Reactions (not covered in the class): (a) Metal Oxide with water Na2O (s) + H2O(l)  2NaOH (aq) (b) Metal Oxide with acid NiO(s) + 2HCl(aq)  NiCl2(aq) + H2O(l)

  23. Supplementary Material • Non-Metals : Tend to gain electrons Usually have negative electron affinity (EA) Seven nonmetals that exist as diatomic molecules H2, N2, O2, F2, Cl2, Br2, I2 Reactions (not covered in the class): (a) Non-Metal Oxide with water P4O10 (s) + 6H2O(l)  4H3PO4 (aq) (b) Non-Metal Oxide with base CO2(g) + 2NaOH(aq)  Na2CO3 (aq) + H2O(l)

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