Chapter 4

What are the wavelength in nanometers and energy in joules of the light emitted when a hydrogen electron originally in the \(n=6\) shell relaxes to the ground state? \(\left[1 \mathrm{eV}=1.602 \times 10^{-19} \mathrm{~J}\right]\)

Use electron configuration notation to explain why aluminum, Al, and gallium, Ga, have similar chemical properties even though gallium has \(d\) electrons but aluminum does not.

Name each element and tell how many valence electrons it has: (a) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{2}\) (b) \([\mathrm{Ne}] 3 \mathrm{~s}^{1}\) (c) \([\mathrm{Ar}] 4 s^{2} 3 d^{10} 4 p^{3}\)

Which subshell is filled in transition metals? What is unique about its order of filling?

Consider the anion whose charge is \(2^{-}\) and whose electron configuration is identical to that of argon, Ar. What is the symbol for this anion?

What do \(\mathrm{F}^{-}, \mathrm{O}^{2-}, \mathrm{Na}^{+}\), and \(\mathrm{Mg}^{2+}\) all have in common?

Why is the formula for calcium sulfide \(\mathrm{CaS}\) and not \(\mathrm{Ca}_{2} \mathrm{~S} ?\)

Select the element from each pair expected to have the lower first ionization energy. Explain. (a) \(\mathrm{N}\) and \(\mathrm{F}\) (b) \(\mathrm{Mg}\) and \(\mathrm{Ba}\) (c) \(\mathrm{N}\) and \(\mathrm{Ca}\)

How does the first ionization energy of the alkali metal in a given period compare with the first ionization energy of the halogen in the same period? What is the result in terms of how these elements tend to react with each other?

Arrange calcium, strontium, arsenic, bromine, and chlorine in order of (a) increasing atomic size and (b) increasing first ionization energy.

Of the atoms \(\mathrm{Na}, \mathrm{Cl}, \mathrm{K}, \mathrm{Br}\), which has the largest atomic radius? Which has the largest first ionization energy?

Arrange \(\mathrm{Ca}, \mathrm{Se}, \mathrm{F}, \mathrm{S}\), and \(\mathrm{Rb}\) in order of increasing ionization energy.

Predict the formula for lithium nitride, made from lithium, \(\mathrm{Li}\), and nitrogen, \(\mathrm{N}\). Explain your reasoning.

Suppose the 3s valence electron in a ground-state sodium atom is excited to the 4 s orbital. Why does the electron now have higher energy?

What is the wavelength in nanometers of electromagnetic waves that have an energy of \(1.00 \times 10^{-15} \mathrm{~J} ?\)

Which electromagnetic radiation is most dangerous to humans? (a) X-rays (b) Ultraviolet light (c) Gamma rays (d) Infrared light (e) Radio waves

Write the ground-state electron configuration for each of the following atoms or ions. Which have a valence-shell octet? (a) \(\mathrm{Ca}^{+}\) (b) Li (c) \(\mathrm{P}^{3-}\) (d) \(\mathrm{Ar}^{+}\) (e) \(\mathrm{Si}^{2+}\)

Halogens are very reactive because (choose the correct answer): (a) They need to gain only one electron to satisfy the octet rule. (b) They have seven electrons in their valence shell, and the more electrons an atom has, the more reactive it is. (c) They are nonmetals, and all nonmetals are reactive. (d) They can easily lose their seven valence electrons to satisfy the octet rule.

Write the ground-state electron configuration for each of the following atoms or ions. Which have a valence-shell octet? (a) \(\overline{\mathrm{Ar}}\) (b) \(\mathrm{Na}^{+}\) (c) \(\mathrm{C}^{2-}\) (d) \(\mathrm{O}^{2-}\) (e) \(\mathrm{Ca}^{2+}\)

Which subshell has the lowest energy? (a) \(4 s\) (b) \(3 p\) (c) \(2 p\) (d) \(3 s\) (e) \(2 s\)

Which is the correct ground-state electron configuration for antimony, Sb? (a) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2} 4 p^{6} 4 d^{10} 5 s^{2} 5 d^{3}\) (b) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2} 4 p^{6} 4 d^{10} 5 s^{2} 5 p^{3}\) (c) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{10} 4 p^{6} 5 s^{2} 4 d^{10} 5 p^{4}\) (d) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{10} 4 p^{6} 5 s^{2} 5 p^{3}\) (e) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{10} 4 p^{6} 5 s^{2} 4 f^{3}\)

What is the energy in joules of green light that has a wavelength of \(500 \mathrm{~nm} ?\)

Indicate whether or not the following volumes are quantized: (a) The volume of water available in \(16-\mathrm{oz}\) bottles (b) The volume of water available from a drinking fountain (c) The volume of soft drink available from a soda-fountain dispenser (d) The volume of soft drink available in \(12-\mathrm{oz}\) cans

Predict the formula of the compound formed by the reaction between: (a) \(\mathrm{Ca}\) and \(\mathrm{Br}\) (b) \(\mathrm{K}\) and \(\mathrm{N}\) (c) \(\mathrm{Al}\) and \(\mathrm{S}\) (d) \(\mathrm{Na}\) and \(\mathrm{I}\) (e) \(\mathrm{Mg}\) and \(\mathrm{O}\)

Arrange in order of increasing valence-shell size: \(\mathrm{Sr}, \mathrm{Mg}, \mathrm{Ba}, \mathrm{Ca}, \mathrm{Be}\).

The order of the colors in the visible spectrum can be remembered by the acronym ROY G. BIV, which stands for red, orange, yellow, green, blue, indigo, violet. In the acronym, are the colors arranged in order of increasing or decreasing wavelength? In order of increasing or decreasing energy?

Write the ground-state electron configuration, both full notation and noble gas abbreviation, and indicate the number of valence electrons for: (a) \(\mathrm{Al}\) (b) I (c) \(\mathrm{Rb}\) (d) \(\overline{\mathrm{Ar}}\) (e) \(\mathrm{Mg}\)

Metals tend to \(\quad\) valence electrons, whereas nonmetals tend to valence electrons.

What is the valence shell in \(\mathrm{Mg}\), Ge, \(\mathrm{W}\), \(C \mathrm{~s} ?\)

If gamma radiation has a wavelength of \(1.00 \times 10^{-12} \mathrm{~m}\), what is the energy of gamma radiation in joules?

What is the formula for the maximum number of electrons in each shell of the Bohr atom? How many electrons are allowed in the \(n=2\) shell? The \(n=6\) shell?

What are the group number, period number, and name of the element whose electron configuration is \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{10} 4 p^{2} ?\)

Circle the correct choice to indicate how many electrons each element must gain or lose to form an octet: \(\begin{array}{llll}\text { (a) } \mathrm{Mg} & \text { gains, loses } & 1,2,3 & \text { electrons }\end{array}\) \(\begin{array}{lll}\text { (b) Se gains, loses } 1,2,3 & \text { electrons }\end{array}\) (c) Al gains, loses \(1,2,3\) electrons \(\begin{array}{llll}\text { (d) Sr gains, loses } & 1,2,3 & \text { electrons }\end{array}\) (e) Br gains, loses \(1,2,3\) electrons (f) P gains, loses \(1,2,3\) electrons

State the Heisenberg uncertainty principle and what it implies about the structure of an atom.

In the Bohr model of the atom, are the electrons in shells closer to the nucleus higher or lower in energy than electrons in shells farther from the nucleus? Explain.

What is the total number of \(p\) -subshell electrons for each of the following atoms: \(\mathrm{P}, \mathrm{Mg}\), Se, \(\mathrm{Zn} ?\)

Write the full ground-state electron configuration for \(\mathrm{Ca}^{2+}, \mathrm{S}^{2-}, \mathrm{Ar}, \mathrm{K}^{+}\)

Rank visible light, gamma rays, X-rays, radio/ television waves, infrared radiation, and ultraviolet light in order of (a) increasing wavelength and (b) increasing energy.

Draw a simple Bohr model (no subshells) for an oxygen atom. How many electrons are in the valence shell? How many more electrons can be put into the valence shell?

Taking all configurations to be for neutral atoms, identify the elements having the following electron configurations: (a) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2} 4 p^{5}\) (b) \(1 s^{2} 2 s^{2} 2 p^{1}\) (c) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{10} 4 p^{6} 5 s^{1}\) (d) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{10} 4 p^{6} 5 s^{2} 4 d^{10} 5 p^{6} 6 s^{2} 4 f^{14} 5 d^{10} 6 p^{6}\) \(7 s^{2} 6 d^{1} 5 f^{3}\) (e) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2}\)

Explain the major difference between the orbits in the Bohr model of the atom and the orbitals in the quantum mechanical model of the atom.

What is the highest value of \(n\) for the electrons in these atoms? (a) \(\overline{\mathrm{Co}}\) (b) As (c) \(\overline{\mathrm{Sr}}\) (d) \(\overline{\mathrm{Po}}\)

The second ionization energy of an atom is the minimum energy required to remove an electron from the \(+1\) cation of the atom, and it is always larger than the first ionization energy. Why is this so? (Hint: Think of atomic electrons as clouds, and each electron must "look" through every other electron cloud in the atom to "see" (feel) the nucleus.)

What is the implication of thinking about atomic electrons as clouds that got Einstein so upset? What did Schrödinger call these clouds?

Predict the formula for the compound aluminum nitride made from the elements aluminum and nitrogen, and explain how you made your prediction.

Consider a \(1 \mathrm{~s} \rightarrow 2 \mathrm{~s}\) electron transition. For which atom would this require shorter wavelength light, \(\mathrm{H}\) or He? Justify your answer.

Which of the following electron transitions in hydrogen would absorb the largest amount of energy? A. \(n=2\) to \(n=3\) B. \(\quad n=2\) to \(n=4\) C. \(n=1\) to \(n=4\) D. \(n=3\) to \(n=1\) E. \(n=7\) to \(n=1\)

True or false: Photons of green light have a higher frequency than photons of blue light. True or false: Photons of blue light have a longer wavelength than photons of orange light.

Arrange the following sets of atoms in order of increasing atomic size: (a) \(\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I}\) (b) \(\mathrm{Mg}\), \(\mathrm{Na}, \mathrm{S}, \mathrm{Al}\) (c) \(\mathrm{Al}, \mathrm{Si}, \mathrm{Ge}, \mathrm{Te}\)

Of the following atoms, which has the largest first ionization energy? (a) \(\mathrm{Br}\) (b) \(\overline{\mathrm{O}}\) (c) \(\mathrm{C}\) (d) \(\mathrm{P}\) (e) I