Chapter 9

Write the electron configuration for each ion. a. \(\mathrm{Cl}^{-}\) b. \(\mathrm{P}^{3-}\) c. \(\mathrm{K}^{+}\) d. \(\mathrm{Mo}^{3+}\) e. \(V^{3+}\)

Write orbital diagrams for each ion and indicate whether the ion is diamagnetic or paramagnetic. a. \(\mathrm{V}^{5+}\) b. \(\mathrm{Cr}^{3+}\) c. \(\mathrm{Ni}^{2+}\) d. \(\mathrm{Fe}^{3+}\)

Write orbital diagrams for each ion and indicate whether the ion is diamagnetic or paramagnetic. a. \(\mathrm{Cd}^{2+}\) b. \(\mathrm{Au}^{+}\) c. \(\mathrm{Mo}^{3+}\) d. \(\mathrm{Zr}^{2+}\)

Which is the larger species in each pair? a. Li or \(L i^{+}\) b. \(\mathrm{I}^{-}\) or \(\mathrm{Cs}^{+}\) c. Cr or \(\mathrm{Cr}^{3+}\) d. \(O\) or \(O^{2-}\)

Which is the larger species in each pair? a. \(\mathrm{Sr}\) or \(\mathrm{Sr}^{2+}\) b. \(\mathrm{N}\) or \(\mathrm{N}^{3-}\) c. \(\mathrm{Ni}\) or \(\mathrm{Ni}^{2+}\) d. \(S^{2-}\) or \(C a^{2+}\)

Arrange this isoelectronic series in order of decreasing radius: \(\mathrm{F}^{-}, \mathrm{O}^{2-}, \mathrm{Mg}^{2+}, \mathrm{Na}^{+} .\)

Arrange this isoelectronic series in order of increasing atomic radius: \(\mathrm{Se}^{2-}, \mathrm{Sr}^{2+}, \mathrm{Rb}^{+}, \mathrm{Br}^{-}\).

Choose the element with the higher first ionization energy from each pair. a. Por I b. \(\mathrm{Si}\) or \(\mathrm{Cl}\) c. \(\mathrm{P}\) or \(\mathrm{Sb}\) d. Ga or Ge

Arrange these elements in order of increasing first ionization energy: \(\mathrm{Si}, \mathrm{F}, \mathrm{In}, \mathrm{N} .\)

Arrange these elements in order of decreasing first ionization energy: \(\mathrm{Cl}, \mathrm{S}, \mathrm{Sn}, \mathrm{Pb}\).

For each element, predict where the "jump" occurs for successive ionization energies. (For example, does the jump occur between the first and second ionization energies, the second and third, or the third and fourth?) a. Be b. \(\mathrm{N}\) c. O d. Li

Choose the element with the more negative (more exothermic) electron affinity from each pair. M b. B or S c. Cor N d. Li or F

Choose the element with the more negative (more exothermic) electron affinity from each pair. a. Mg or S b. \(\mathrm{K}\) or \(\mathrm{Cs}\) c. Si or P d. Ga or Br

Choose the more metallic element from each pair. a. Sr or Sb b. As or Bi c. Cl or O d. S or As

Choose the element with the more negative (more exothermic) electron affinity from each pair. a. \(\mathrm{Mg}\) or \(\mathrm{S}\) b. \(\mathrm{Kor} \mathrm{Cs}\) c. \(\mathrm{Si}\) or \(\mathrm{P}\) d. Ga or Br

Choose the more metallic element from each pair. a. Sb or \(\mathrm{Pb}\) b. K or Ge c. Ge or \(S b\) d. As or \(\mathrm{Sn}\)

Arrange these elements in order of decreasing metallic character: \(\mathrm{Sr}, \mathrm{N}, \mathrm{Si}, \mathrm{P}, \mathrm{Ga}, \mathrm{Al}\).

Bromine is a highly reactive liquid while krypton is an inert gas. Explain this difference based on their electron configurations.

Potassium is a highly reactive metal while argon is an inert gas. Explain this difference based on their electron configurations.

Both vanadium and its \(3+\) ion are paramagnetic. Refer to their electron configurations to explain this statement.

Suppose you were trying to find a substitute for \(\mathrm{Na}^{+}\) in nerve signal transmission. Where would you begin your search? What ions would be most like \(\mathrm{Na}^{+} ?\) For each ion you propose, explain the ways in which it would be similar to \(\mathrm{Na}^{+}\) and the ways it would be different. Use periodic trends in your discussion.

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?

Which pair of elements would you expect to have the most similar atomic radii, and why? a. Si and Ga b. Si and Ge c. Si and As

Consider these elements: \(\mathrm{N}, \mathrm{Mg}, \mathrm{O}, \mathrm{F}, \mathrm{Al} .\) 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 \(\mathrm{b}\) and \(\mathrm{c} .\)

Consider these elements: \(\mathrm{P}, \mathrm{Ca}, \mathrm{Si}, \mathrm{S}, \mathrm{Ga}\) 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.

Explain why atomic radius decreases as you move to the right across a period for main-group elements but not for transition elements.

The lightest noble gases, such as helium and neon, are completely inert-they do not form any chemical compounds whatsoever. The heavier noble gases, in contrast, do form a limited number of compounds. Explain this difference in terms of trends in fundamental periodic properties.

The lightest halogen is also the most chemically reactive, and reactivity generally decreases as you move down the column of halogens in the periodic table. Explain this trend in terms of periodic properties.

Write general outer electron configurations \(\left(n s^{x} n p^{y}\right)\) for groups \(6 \mathrm{~A}\) and \(7 \mathrm{~A}\) in the periodic table. The electron affinity of each group 7A element is more negative than that of each corresponding group 6A element. Use the electron configurations to explain why this is so.

The electron affinity of each group 5 A element is more positive than that of each corresponding group 4 A element. Use the outer electron configurations for these columns to suggest a reason for this observation.

The elements with atomic numbers 35 and 53 have similar chemical properties. Based on their electronic configurations, predict the atomic number of a heavier element that also should share these chemical properties.

The electron affinity of sodium is lower than that of lithium, while the electron affinity of chlorine is higher than that of fluorine. Suggest an explanation for this observation.

Use Coulomb's law to calculate the ionization energy in \(\mathrm{kJ} / \mathrm{mol}\) of an atom composed of a proton and an electron separated by \(100.00 \mathrm{pm} .\) What wavelength of light has sufficient energy to ionize the atom?

Consider the elements: \(\mathrm{B}, \mathrm{C}, \mathrm{N}, \mathrm{O}, \mathrm{F}\) a. Which element has the highest first ionization energy? b. Which element has the largest atomic radius? c. Which element is most metallic? d. Which element has three unpaired electrons? Consider the elements: \(\mathrm{Na}, \mathrm{Mg}, \mathrm{Al}, \mathrm{Si}, \mathrm{P}\)

Consider the elements: \(\mathrm{Na}, \mathrm{Mg}, \mathrm{Al}, \mathrm{Si}, \mathrm{P}\) a. Which element has the highest second ionization energy? b. Which element has the smallest atomic radius? c. Which element is least metallic? d. Which element is diamagnetic?

Consider the metals in the first transition series. Use periodic trends to predict a trend in density as you move to the right across the series.

Only trace amounts of the synthetic element darmstadtium, atomic number 110 , have been obtained. The element is so highly unstable that no observations of its properties have been possible. Based on its position in the periodic table, propose three different reasonable valence electron configurations for this element.

The trend in second ionization energy for the elements from lithium to fluorine is not a regular one. Predict which of these elements has the highest second ionization energy and which has the lowest and explain. Of the elements \(\mathrm{N}, \mathrm{O},\) and \(\mathrm{F}, \mathrm{O}\) has the highest and \(\mathrm{N}\) the lowest second ionization energy. Explain.

Even though adding two electrons to \(\mathrm{O}\) or \(\mathrm{S}\) forms an ion with a noble gas electron configuration, the second electron affinity of both of these elements is positive. Explain.

The outermost valence electron in atom A experiences an effective nuclear charge of \(2+\) and is on average \(225 \mathrm{pm}\) from the nucleus. The outermost valence electron in atom B experiences an effective nuclear charge of \(1+\) and is on average \(175 \mathrm{pm}\) from the nucleus. Which atom (A or B) has the higher first ionization energy? Explain.

Determine whether each statement regarding penetration and shielding is true or false. (Assume that all lower energy orbitals are fully occupied.) a. An electron in a \(3 s\) orbital is more shielded than an electron in a \(2 s\) orbital. b. An electron in a \(3 s\) orbital penetrates into the region occupied by core electrons more than electrons in a \(3 p\) orbital penetrates into the region occupied by core electrons. c. An electron in an orbital that penetrates closer to the nucleus always experiences more shielding than an electron in an orbital that does not penetrate as far. d. An electron in an orbital that penetrates close to the nucleus tends to experience a higher effective nuclear charge than an electron in an orbital that does not penetrate close to the nucleus.

Give a combination of four quantum numbers that could be assigned to an electron occupying a \(5 p\) orbital. Do the same for an electron occupying a \(6 d\) orbital.

Use the trends in ionization energy and electron affinity to explain why calcium fluoride has the formula \(\mathrm{CaF}_{2}\) and not \(\mathrm{Ca}_{2} \mathrm{~F}\) or CaF.

Have each member of your group sketch a periodic table indicating a periodic trend (atomic size, first ionization energy, metallic character, etc.). Have each member present his or her table to the rest of the group and explain the trend based on concepts such as orbital size or effective nuclear charge.