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Chapter 4: Problem 148

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?

Short Answer

Expert verified
In a given period, the first ionization energy of an alkali metal is lower than that of a halogen. Alkali metals have a single valence electron, which is easily removed due to its distance from the nucleus, leading to low ionization energy. Halogens, on the other hand, have higher ionization energies because they are one electron short of a full outer energy level, and their valence electrons are closer to the nucleus. This difference in ionization energies makes alkali metals and halogens highly reactive with each other. When they react, alkali metals lose their valence electron to halogens, forming ionic compounds known as alkali metal halides, such as NaCl and LiF.

Step by step solution

01

Define Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The first ionization energy is the energy needed to remove the first electron from a gaseous atom. Generally, ionization energies increase as we move across a period from left to right in the periodic table, and decrease as we move down a group.
02

Determine Ionization Energy Trends for Alkali Metals and Halogens

Alkali metals (Group 1 elements) have the lowest ionization energies in their respective periods. This is because they have only one electron in their outermost energy level (valence electron), which is relatively easy to remove due to its distance from the nucleus. Halogens (Group 17 elements) have higher ionization energies compared to alkali metals in the same period. This is due to their electron configuration, which is only one electron short of a full outer energy level. The longer electron-nucleus distance for the alkali metal's valence electron makes it easier to remove than the halogen's valence electron, which is closer to the nucleus.
03

Compare Ionization Energies of Alkali Metals and Halogens

In each period, the first ionization energy of an alkali metal is lower than the first ionization energy of a halogen. For example, in period 2, lithium (an alkali metal) has a first ionization energy of 520 kJ/mol, while fluorine (a halogen) has a first ionization energy of 1681 kJ/mol. This trend holds for each period in the periodic table, as alkali metals consistently have lower first ionization energies than halogens.
04

Analyze Reactivity Based on Ionization Energies

The difference in ionization energies between alkali metals and halogens makes them highly reactive with each other. Alkali metals, having a low ionization energy, easily lose the outermost electron to form positively charged ions (cations). On the other hand, halogens, with high ionization energy, readily gain an electron to achieve a full outer energy level, forming negatively charged ions (anions).
05

Predict Reaction Between Alkali Metals and Halogens

When alkali metals and halogens react, they tend to form ionic compounds. In these reactions, the alkali metal loses its valence electron to a halogen, creating a stable ionic bond between the resulting cation (alkali metal) and anion (halogen). These ionic bonds result in the formation of alkali metal halides, such as NaCl (sodium chloride) or LiF (lithium fluoride), which are often found as salts in nature.

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Most popular questions from this chapter

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

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.

Below are data on the first four ionization energies for a fictitious element \(X\). First ionization energy \(=500 \mathrm{~kJ} / \mathrm{mol}\) Second ionization energy \(=2000 \mathrm{~kJ} / \mathrm{mol}\) Third ionization energy \(=3500 \mathrm{~kJ} / \mathrm{mol}\) Fourth ionization energy \(=25,000 \mathrm{~kJ} / \mathrm{mol}\) From the data, which of the following statements is(are) incorrect? (a) \(\mathrm{X}\) could belong to group IIIA. (b) The fourth ionization energy is so much greater than the third ionization energy because \(\mathrm{X}^{3+}\) consists of a noble-gas core or a pseudo- noble-gas core. (c) The third ionization energy is greater than the second ionization energy because \(\mathrm{X}^{2+}\) has a bigger charge than \(\mathrm{X}^{+}\). (d) \(\mathrm{X}\) could belong to group IIIB. (e) \(X\) could belong to group VA.

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