# 5.04 – Periodic Trends

When we group matter together, we find that certain elements share properties.  Many of these properties have a direct connection with the electronic structure of the atom itself.  In this mission, we will go into detail on some of these trends, to help us make predictions using the periodic table.

Lesson Objectives

• Explain the meaning of effective nuclear charge, Zeff, and how Zeff depends on nuclear charge and electron configuration. (7.2)
• Predict the trends in atomic radii, ionic radii, ionization energy, and electron affinity by using the periodic table. (7.2, 7.3, 7.4, and 7.5)
• Explain how the radius of an atom changes upon losing electrons to form a cation or gaining electrons to form an anion. (7.3)
• Write the electron configurations of ions. (7.3)
• Explain how the ionization energy changes as we remove successive electrons, and the jump in ionization energy that occurs when the ionization corresponds to removing a core electron. (7.4)
• Explain how irregularities in the periodic trends for electron affinity can be related to electron configuration. (7.5)

As we begin this topic, make sure to skim the text and take down notes.  The key ideas are presented in the slides below, which correspond to sections 7.2-4.5 of the text.

Now that you've seen the content, work out the sample and practice problems below (sample problems 7.1 - 7.7 of the text).

To wrap up this mission, work out the mastery problems below on a separate sheet of paper, then turn in all of your stamped work on OneNote.  These problems come from the back of chapter 7 in your book.

7.16    Arrange the following atoms in order of increasing effective nuclear charge experienced by the electrons in the n = 3 electron shell: K, Mg, P, Rh, Ti.  Explain the basis for your order.

7.35    Consider S, Cl, and K and their most common ions.

(a) List the atoms in order of increasing size.

(b) List the ions in order of increasing size.

(c) Explain any differences in the orders of the atomic and ionic sizes.

7.97    Use electron configurations to explain the following observations:

(a) The first ionization energy of phosphorus is greater than that of sulfur.

(b) The electron affinity of nitrogen is lower (less negative) than those of both carbon and oxygen.

(c) The second ionization energy of oxygen is greater than the first ionization energy of fluorine.

(d) The third ionization energy of manganese is greater than those of both chromium and iron.

7.107  One way to measure ionization energies is ultraviolet photoelectron spectroscopy (PES), a technique based on the photoelectric effect. (Section 6.2) In PES, monochromatic light is directed onto a sample, causing electrons to be emitted. The kinetic energy of the emitted electrons is measured. The difference between the energy of the photons and the kinetic energy of the electrons corresponds to the energy needed to remove the electrons (that is, the ionization energy). Suppose that a PES experiment is performed in which mercury vapor is irradiated with ultraviolet light of wavelength 58.4 nm.

(a) What is the energy of a photon of this light, in eV?

(b) Write an equation that shows the process corresponding to the first ionization energy of Hg.

(c) The kinetic energy of the emitted electrons is measured to be 10.75 eV. What is the first ionization energy of Hg, in eV

(d) Using Figure 7.10, determine which of the halogen elements has a first ionization energy closest to that of mercury.

2003b FRQ #7

Well Visitor, you're certainly going out of you way to stay on top of things.  Great work :)