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Chapter 8: Problem 67

Use the alkali metals and alkaline earth metals as examples to show how we can predict the chemical properties of elements simply from their electron configurations.

Short Answer

Expert verified
By knowing the electron configurations of alkali metals (\(s^1\)) and alkaline earth metals (\(s^2\)), we know these metals tend to lose these valence electrons to achieve a stable electron configuration, as seen in their chemical reactivity, forming \(+1\) and \(+2\) ions respectively. Also since elements in same group have similar configurations, they exhibit similar chemical properties.

Step by step solution

01

Understanding electron configuration

Alkali metals are in Group 1 and have an electron configuration that ends in \(s^1\). For example, Sodium (Na) has electron configuration \([Ne]3s^1\). Alkaline earth metals are in Group 2 and the electron configuration ends in \(s^2\). For example, Magnesium (Mg) with electron configuration \([Ne]3s^2\).
02

Link electron configuration to reactivity

These electron configurations particularly the outermost electrons (valence electrons) determine the reactivity of the elements. Alkali metals have 1 valence electron which they tend to lose to achieve a stable electron configuration. They thus often form \(+1\) ions in chemical reactions. Similarly, alkaline earth metals have 2 valence electrons which they tend to lose and thus form \(+2\) ions.
03

Predict physical properties

This understanding of electron configurations and valence electrons can help us predict the chemical properties. Alkali and alkaline earth metals are very reactive metals. Alkali metals are more reactive than alkaline earth metals due to fewer valence electrons making it easier to lose that single electron and achieve stability.
04

Further generalization

Moreover, elements in the same group of the periodic table have similar electron configurations (same number of valence electrons) and therefore show similar chemical properties. Thus we can predict the chemical properties of elements within groups of the periodic table using their electron configurations.

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Most popular questions from this chapter

Define ionic radius. How does the size of an atom change when it is converted to (a) an anion and (b) a cation?

An element X reacts with hydrogen gas at \(200^{\circ} \mathrm{C}\) to form compound Y. When \(Y\) is heated to a higher temperature, it decomposes to the element \(X\) and hydrogen gas in the ratio of \(559 \mathrm{~mL}\) of \(\mathrm{H}_{2}\) (measured at STP) for \(1.00 \mathrm{~g}\) of \(\mathrm{X}\) reacted. \(\mathrm{X}\) also combines with chlorine to form a compound \(Z,\) which contains 63.89 percent by mass of chlorine. Deduce the identity of \(X\).

As a group, the noble gases are very stable chemically (only \(\mathrm{Kr}\) and Xe are known to form compounds). Use the concepts of shielding and the effective nuclear charge to explain why the noble gases tend to neither give up electrons nor accept additional electrons.

To prevent the formation of oxides, peroxides, and superoxides, alkali metals are sometimes stored in an inert atmosphere. Which of the following gases should not be used for lithium: \(\mathrm{Ne}, \mathrm{Ar}, \mathrm{N}_{2}, \mathrm{Kr} ?\) Explain. (Hint: As mentioned in the chapter, Li and Mg exhibit a diagonal relationship. Compare the common compounds of these two elements.)

(a) Define electron affinity. (b) Electron affinity measurements are made with gaseous atoms. Why? (c) Ionization energy is always a positive quantity, whereas electron affinity may be either positive or negative. Explain.
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