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Chapter 8: Problem 31

Compare the elements Na, Mg, O, and P. (a) Which has the largest atomic radius? (b) Which has the most negative electron affinity? (c) Place the elements in order of increasing ionization energy.

Short Answer

Expert verified
(a) Na; (b) O; (c) Na < Mg < P < O.

Step by step solution

01

Identify Atomic Radius Trend

Atomic radius generally decreases across a period (from left to right) on the periodic table and increases down a group. Sodium (Na) is located furthest to the left and therefore has the largest atomic radius among Na, Mg, O, and P.
02

Determine Electron Affinity Characteristics

Electron affinity becomes more negative as you move across a period to the right (excluding the noble gases). Between the given elements, oxygen (O) has the most negative electron affinity because it is to the right of Na, Mg, and P on the periodic table.
03

Assess Ionization Energy Trend

Ionization energy increases across a period from left to right and decreases down a group. Thus, the order of increasing ionization energy from lowest to highest for the elements Na, Mg, O, and P is Na < Mg < P < O.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Atomic Radius
Atomic radius is a fundamental concept when exploring periodic trends in elements. It refers to the size of an atom from its nucleus to the outermost electron shell. In the periodic table, two major trends affect atomic radius:
  • Across a period (left to right), the atomic radius decreases. This happens because more protons are added to the nucleus, attracting the electrons more strongly and pulling them closer.
  • Down a group (top to bottom), the atomic radius increases. As we go down a group, a new electron shell is added, further increasing the distance between the nucleus and the outermost electrons.
For elements like sodium (Na), magnesium (Mg), oxygen (O), and phosphorus (P), sodium, being furthest to the left in its period, has the largest atomic radius. This is because it has fewer protons compared to the others, leading to a weaker overall pull on its electrons.
Electron Affinity
Electron affinity is a measure of how much energy is released when an atom in the gaseous state gains an electron. It gives us insight into how eager an atom is to accept an additional electron. When considering periodic trends:
  • Across a period, electron affinity tends to become more negative, reflecting an increase in the atom's affinity for an electron.
  • Down a group, electron affinity usually becomes less negative, indicating a lesser tendency for atoms to gain electrons.
For the elements sodium (Na), magnesium (Mg), oxygen (O), and phosphorus (P), oxygen becomes significant. Oxygen, being further to the right in the period, has a very negative electron affinity due to a higher attraction for electrons--more than Na, Mg, or P. This reflects its strong desire to fill its outer electron shell to achieve a stable electronic configuration.
Ionization Energy
Ionization energy is the energy required to remove an electron from an atom in its gaseous state. It reflects an atom's hold on its electrons and its capacity to avoid losing them. Key periodic trends for ionization energy include:
  • Across a period, ionization energy generally increases. This is because adding more protons to the nucleus increases nuclear charge, enhancing the nucleus's attraction to electrons.
  • Down a group, ionization energy decreases. This is due to the increased distance between the nucleus and outer electrons, along with greater electron shielding, reducing the pull on the outermost electrons.
Considering elements like sodium (Na), magnesium (Mg), oxygen (O), and phosphorus (P), the order of increasing ionization energy is Na < Mg < P < O. Oxygen, having the highest ionization energy, reflects the strongest hold on its electrons--more than the other three elements.

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Most popular questions from this chapter

Name the element corresponding to each characteristic below. (a) the element with the electron configuration \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{3}\) (b) the alkaline earth element with the smallest atomic radius (c) the element with the largest ionization energy in Group \(5 \mathrm{A}\) (d) the element whose \(2+\) ion has the configuration \([\mathrm{Kr}] 4 d^{5}\) (e) the element with the most negative electron affinity in Group 7A (f) the element whose electron configuration is \([\mathrm{Ar}] 3 d^{10} 4 s^{2}\)

Using orbital box diagrams, depict an electron configuration for each of the following ions: (a) \(\mathrm{Na}^{+},\) (b) \(\mathrm{Al}^{3+}\) (c) \(\mathrm{Ge}^{2+},\) and \((\mathrm{d}) \mathrm{F}^{-}\).

Explain briefly why each of the following is not a possible set of quantum numbers for an electron in an atom. In each case, change the incorrect value (or values) to make the set valid. (a) \(n=2, \ell=2, m_{\ell}=0, m_{\mathrm{s}}=+\frac{1}{2}\) (b) \(n=2, \ell=1, m_{\ell}=-1, m_{\mathrm{s}}=0\) (c) \(n=3, \ell=1, m_{\ell}=+2, m_{\mathrm{s}}=+\frac{1}{2}\)

What is the maximum number of electrons that can be identified with each of the following sets of quantum numbers? In some cases, the answer may be "none." In such cases, explain why "none" is the correct answer. (a) \(n=3\) (b) \(n=3\) and \(\ell=2\) (c) \(n=4, \ell=1, m_{\ell}=-1,\) and \(m_{\mathrm{s}}=-\frac{1}{2}\) (d) \(n=5, \ell=0, m_{\ell}=+1\)

Using an orbital box diagram and noble gas notation, show the electron configurations of uranium and of the uranium(IV) ion. Is either of these paramagnetic?
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