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Refer to their electron configurations to explain why copper is paramagnetic, whereas its 1+ ion is not.
Ch.8 - Periodic Properties of the Elements
Chapter 8, Problem 101
Explain why atomic radius decreases as you move to the right across a period for main-group elements but not for transition elements.
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Understand that atomic radius is the distance from the nucleus of an atom to the outermost electron shell.
Recognize that as you move to the right across a period in the periodic table, the number of protons in the nucleus increases, which increases the positive charge of the nucleus.
Acknowledge that the increased nuclear charge attracts the electrons more strongly, pulling them closer to the nucleus and resulting in a smaller atomic radius for main-group elements.
Note that for transition elements, the additional electrons are added to an inner d-subshell rather than the outermost shell, which does not significantly increase the shielding effect.
Understand that the effective nuclear charge experienced by the outer electrons in transition elements does not increase as much as in main-group elements, so the atomic radius does not decrease significantly across a period for transition elements.
Related Practice
Textbook Question
Life on Earth evolved based on the element carbon. Based on periodic properties, what two or three elements would you expect to be most like carbon?
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Consider these elements: N, Mg, O, F, Al. a. Write the electron configuration for each element. b. Arrange the elements in order of decreasing atomic radius. c. Arrange the elements in order of increasing ionization energy. d. Use the electron configurations in part a to explain the differences between your answers to parts b and c.
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Explain why vanadium (radius = 134 pm) and copper (radius = 128 pm) have nearly identical atomic radii, even though the atomic number of copper is about 25% higher than that of vanadium. Predict the relative densities of these two metals. Look up the densities in a reference book, periodic table, or on the Internet to check if your predictions are correct.
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The lightest noble gases, such as helium and neon, are completely inert—they do not form any chemical compounds whatsoever. In contrast, the heavier noble gases do form a limited number of compounds. Explain this difference in terms of trends in fundamental periodic properties.
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Why does the lightest halogen, which is also the most chemically reactive, exhibit a decrease in reactivity as you move down the column of halogens in the periodic table? Explain this trend in terms of periodic properties.