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Chemical Periodicity Chart

Chemical Periodicity Chart. Practice Problem Questions and Answers. Atomic Radius. P. 178 #16, 22 16: How does atomic size change within groups and across periods? Increases down the groups, decreases left-to-right across periods.

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Chemical Periodicity Chart

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  1. Chemical Periodicity Chart Practice Problem Questions and Answers

  2. Atomic Radius • P. 178 #16, 22 • 16: How does atomic size change within groups and across periods? • Increases down the groups, decreases left-to-right across periods. • 22: Arrange in order of decreasing size: S, Cl, Al, Na. Is there a pattern? • Na, Al, S, Cl • This is a period-based trend (left-to-right, n=3).

  3. Atomic Radius • P.181 #36 • A: Which element has a larger radius: Na or Li? • Na • B: Which element has a larger radius: Sr or Mg? • Sr • C: Which element has a larger radius: C or Ge? • Ge • D: Which element has a larger radius: O or Se? • Se

  4. Atomic Radius • P. 182 #50 • Why does fluorine have a smaller atomic radius than oxygen and chlorine? • It’s further to the right in oxygen’s period, it’s higher up than chlorine. • In other words, “stronger nucleus” than oxygen, fewer electrons than chlorine.

  5. Ionization Energy • P. 178 #17-18 • 17: When do ions form? • When electrons are added or removed. • 18: What happens to first ionization energy within groups and across periods? • Increases left-to-right across periods, decreases down groups.

  6. Ionization Energy • P.178 #23 • A: Which element has the larger first ionization energy: Na, K? • Na • B: Which element has the larger first ionization energy: Mg, P? • P

  7. Ionization Energy • P.181 #37, 38 • 37: Explain the difference between first and second ionization energy: • First i.e. = energy to remove one electron. • Second i.e. = energy to remove a second electron • 38: Which element has a greater first i.e.? • Li, B • B • Mg, Sr • Mg • Cs, Al • Al

  8. Ionization Energy • P. 181 #39 • Arrange the groups of elements in order of increasing ionization energy: • Be, Mg, Sr • Sr, Mg, Be • Bi, Cs, Ba • Cs, Ba, Bi • Na, Al, S • Na, Al, S

  9. Ionization Energy • P.181 #40 • Why is there a large increase between the first and second ionization energies of the alkali metals? • After removing the first electron, the second electron is in a lower (closer) energy level (lower n number).

  10. Ionization Energy • P. 182: 51, 55 • 51: Would you expect metals or nonmetals in the same period to have higher i.e.? • Nonmetals – they’re further right (“stronger nuclei”) • 55: Which equation represents the first ionization of an alkali metal atom? • A: Cl Cl+ + e- • B: Ca  Ca+ + e- • C: K  K+ + e- • D: H  H+ + e-

  11. Ionization Energy • P. 182: 51, 55 • 51: Would you expect metals or nonmetals in the same period to have higher i.e.? • Nonmetals – they’re further right (“stronger nuclei”) • 55: Which equation represents the first ionization of an alkali metal atom? • A: Cl Cl+ + e- • B: Ca  Ca+ + e- • C: K  K+ + e- • D: H  H+ + e-

  12. Ionization Energy • P.182 #58 • Why is there a large jump between the second and third ionization energies of magnesium? Why is there a large jump between the third and fourth ionization energies of aluminum? • Those last electrons are in closer energy shells (lower n number).

  13. Ionic Size • P. 178 #19 • Compare the size of ions to the size of their neutral forms. • Cations lose electrons, become positively charged, get smaller. • Anions gain electrons, become negatively charged, get larger.

  14. Ionic Size • P. 181 #41, 42 • 41: How does the ionic radius of a typical metal compare with its atomic radius? • Metals tend to lose electrons so their ionic radii get smaller. • 42: Which particle has a larger radius in each atom/ion pair? • Na, Na+ • Na • S, S2- • S2- • I, I- • I- • Al, Al3- • Al

  15. Ionic Size • P. 182: #52 • In each pair, which ion is larger? • Ca2+, Mg2+ • Ca2+ • Cl-, P3- • P3- • Cu+, Cu2+ • Cu+

  16. Ionic Size • P. 182 #59 • The bar graph shows the relationship between atomic and ionic radii for Group 1A elements. A: Describe the trend in atomic radius. B: Explain the difference between ionic and atomic radius size? • A: Radius increases as you go down a group. • B: Ions are smaller due to fewer electrons than in the neutral atom (atomic radius).

  17. Ionic Size • P.183 #64, 65 • 64: The Mg2+ and Na+ ions each have ten electrons. Which is smaller and why? • Mg2+is smaller because though it has ten electrons just like Na+, it has more protons. They pull “harder” on the electrons. • 65: How do you expect the radii of S2-, Cl-, K+, Ca2+, and Sc3+ to vary – they have the same total electrons as the noble gas Argon. What about for O2-, F-, Na+, Mg2+, and Al3+, which is the same as Neon? • Radius decreases from left to right across a period in both cases. Though electron # is the same, proton number goes up.

  18. Ionic Size • P. 183 #68 • Atoms and ions with the same number of electrons are isoelectronic. • Write the symbol for a cation and anion that are isoelectronic with Krypton: • Br-, Rb+, Se2-, As3-, Sr2+(each have 36 electrons) • Can you have an isoelectriccation and anion in the same period? • No, cations lose electrons but anions (higher overall number of electrons) gain them.

  19. Electronegativity • P. 178 #20 • How does electronegativity vary within groups and across periods? • Increases across period left-to-right. • Decreases down groups.

  20. Electronegativity • P.181 #43 • A: Which element has a higher electronegativity value: Cl, F? • F • B: Which element has a higher electronegativity value: C, N? • N • C: Which element has a higher electronegativity value: Mg, Ne? • Mg [Ne does not react] • D: Which element has a higher electronegativity value: As, Ca? • As

  21. Electronegativity • P.181 #44 • Why are noble gases not given electronegativity values? • Electronegativity only applies in compounds. Noble gases don’t react and form compounds.

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