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Chapter 7: Problem 67

(a) Why is calcium generally more reactive than magnesium? (b) Why is calcium generally less reactive than potassium?

Short Answer

Expert verified
Calcium is generally more reactive than magnesium because it has a larger atomic radius, meaning the distance between its nucleus and valence electrons is greater. This weaker attractive force makes it easier for calcium to lose its electrons. On the other hand, calcium is generally less reactive than potassium because potassium is in Group 1 of the periodic table, making it easier for potassium to lose its single valence electron than calcium to lose its two valence electrons. The attractive force between the nucleus and valence electrons is weaker for potassium than for calcium, making it easier for potassium to lose its valence electron and form an ion.

Step by step solution

01

1. Electron configurations

Write down the electron configurations of calcium (Ca), magnesium (Mg), and potassium (K): - Ca (Atomic number 20): \(1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{2}\) - Mg (Atomic number 12): \(1s^{2} 2s^{2} 2p^{6} 3s^{2}\) - K (Atomic number 19): \(1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{1}\)
02

2. Valence electrons and reactivity

The reactivity of elements increases with the ease of losing valence electrons or gaining valence electrons. Generally, metals (like Ca, Mg, and K) tend to lose electrons to form positive ions, while nonmetals gain electrons to form negative ions. The number of valence electrons and the energy needed to remove them affect the reactivity of a metal. For the given elements, their valence electrons and ease of removal are as follows: - Ca has 2 valence electrons in the fourth energy level (4s) - Mg has 2 valence electrons in the third energy level (3s) - K has 1 valence electron in the fourth energy level (4s)
03

3. Reactivity trend in the periodic table

As we move down a group in the periodic table, the distance between the nucleus and the valence electrons increases. This leads to a weaker attractive force between the nucleus and valence electrons, making it easier for atoms to lose electrons, subsequently increasing their reactivity. However, as we move across a period from left to right in the periodic table, the number of protons in the nucleus increases, while the valence electrons remain in the same energy level, resulting in a stronger attractive force and making it harder for these elements to lose electrons. Therefore, their reactivity decreases.
04

4. Reactivity comparison: Calcium vs. Magnesium

Calcium (Ca) and magnesium (Mg) are both in Group 2 of the periodic table, with Ca located below Mg. Since Ca has a larger atomic radius, the distance between its nucleus and valence electrons is greater than that of Mg. Consequently, there is a weaker attractive force between the nucleus and valence electrons in Ca, making it easier to lose electrons. This is the reason why calcium is generally more reactive than magnesium.
05

5. Reactivity comparison: Calcium vs. Potassium

Calcium (Ca) and potassium (K) are in different groups of the periodic table, with K being in Group 1 and Ca in Group 2. Both elements are in the same period, but K is one group to the left of Ca, making it easier for K to lose its single valence electron than it is for Ca to lose its two valence electrons. The attractive force between the nucleus and valence electrons is weaker for K than for Ca, making it easier for K to lose its valence electron and forming an ion. Consequently, this is why calcium is generally less reactive than potassium.

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Most popular questions from this chapter

Mercury in the environment can exist in oxidation states 0, +1, and +2. One major question in environmental chemistry research is how to best measure the oxidation state of mercury in natural systems; this is made more complicated by the fact that mercury can be reduced or oxidized on surfaces differently than it would be if it were free in solution. XPS, X-ray photoelectron spectroscopy, is a technique related to PES (see Exercise 7.111), but instead of using ultraviolet light to eject valence electrons, X rays are used to eject core electrons. The energies of the core electrons are different for different oxidation states of the element. In one set of experiments, researchers examined mercury contamination of minerals in water. They measured the XPS signals that corresponded to electrons ejected from mercury’s 4\(f\) orbitals at 105 eV, from an X-ray source that provided 1253.6 \(\mathrm{eV}\) of energy \(\left(1 \mathrm{ev}=1.602 \times 10^{-19} \mathrm{J}\right)\) The oxygen on the mineral surface gave emitted electron energies at \(531 \mathrm{eV},\) corresponding to the 1 \(\mathrm{s}\) orbital of oxygen. Overall the researchers concluded that oxidation states were \(+2\) for \(\mathrm{Hg}\) and \(-2\) for \(\mathrm{O}\) (a) Calculate the wavelength of the X rays used in this experiment. (b) Compare the energies of the 4f electrons in mercury and the 1s electrons in oxygen from these data to the first ionization energies of mercury and oxygen from the data in this chapter. (c) Write out the ground- state electron configurations for \(\mathrm{Hg}^{2+}\) and \(\mathrm{O}^{2-} ;\) which electrons are the valence electrons in each case?

Hydrogen is an unusual element because it behaves in some ways like the alkali metal elements and in other ways like nonmetals. Its properties can be explained in part by its electron configuration and by the values for its ionization energy and electron affinity. (a) Explain why the electron affinity of hydrogen is much closer to the values for the alkali elements than for the halogens. (b) Is the following statement true? "Hydrogen has the smallest bonding atomic radius of any element that forms chemical compounds." If not, correct it. If it is, explain in terms of electron configurations. (c) Explain why the ionization energy of hydrogen is closer to the values for the halogens than for the alkali metals. (d) The hydride ion is \(\mathrm{H}^{-} .\) Write out the process corresponding to the first ionization energy of the hydride ion. (e) How does the process in part (d) compare to the process for the electron affinity of a neutral hydrogen atom?

Give three examples of ions that have an electron configuration of \(n d^{6}(n=3,4,5, \ldots).\)

The electron affinities, in \(\mathrm{kJ} / \mathrm{mol},\) for the group 1 \(\mathrm{B}\) and group 2 \(\mathrm{B}\) metals are as follows: (a) Why are the electron affinities of the group 2 \(\mathrm{B}\) elements greater than zero? (b) Why do the electron affinities of the group 1 \(\mathrm{B}\) elements become more negative as we move down the group? [Hint: Examine the trends in the electron affinities of other groups as we proceed down the periodic table.]

Write balanced equations for the following reactions: (a) barium oxide with water, (b) iron(II) oxide with perchloric acid, (c) sulfur trioxide with water, (d) carbon dioxide with aqueous sodium hydroxide.
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