Question

Which of the following statements is(are) true? a. F has a larger first ionization energy than does Li. b. Cations are larger than their parent atoms. c. The removal of the first electron from a lithium atom (electron configuration is 1\(s^{2} 2 s^{1} )\) is exothermic - that is, removing this electron gives off energy. d. The He atom is larger than the \(\mathrm{H}^{+}\) ion. e. The Al atom is smaller than the Li atom.

Short answer

Statements a and e are true. Statement a is true because F has a larger first ionization energy than Li due to its position in the periodic table. Statement e is true because the Al atom is smaller than the Li atom due to the increased number of electron shells and a higher effective nuclear charge. Statements b, c, and d are false.

Step-by-step solution

Statement a: Comparing First Ionization Energies of F and Li

First ionization energy is the energy required to remove the outermost electron from a neutral atom. As we move across a period (row) in the periodic table from left to right, the ionization energy generally increases because of the increase in effective nuclear charge experienced by electrons, as well as the decrease in atomic size. F (Fluorine) is in the same period as Li (Lithium), but to the right of it. Hence, F's first ionization energy should be larger than Li's. So, statement a is true.

Statement b: Comparing Cations to their Parent Atoms

When an atom loses one or more electrons, it forms cations. Since electrons are negatively charged, the loss of electrons reduces the repulsion between the electrons and causes the remaining electrons to be drawn closer to the nucleus. Thus, cations are smaller than their parent atoms. Statement b is false, as it claims the opposite.

Statement c: Removing an Electron from Lithium

To remove the first electron from a lithium atom, which has an electron configuration of 1s² 2s¹, energy needs to be supplied to overcome the attraction between the electron and the nucleus. Hence, the process of removing an electron from an atom is endothermic and requires energy input. Statement c is false because it wrongly claims that the removal of the electron is exothermic.

Statement d: Comparing Size of He Atom and H+ Ion

A helium (He) atom has an electron configuration of 1s², while a hydrogen ion (\(\mathrm{H}^{+}\)) has lost its only electron, and the remaining nucleus consists of a single proton. Since electrons contribute to the size of an atom or ion, the \(\mathrm{H}^{+}\) ion will be smaller than a He atom since it has no electrons. Statement d is false because it claims that the He atom is larger than the \(\mathrm{H}^{+}\) ion.

Statement e: Comparing Sizes of Al and Li Atoms

As we move down a group (column) in the periodic table, the size of atoms generally increases due to an increase in the number of electron shells. However, as we move across a period from left to right, the size of atoms generally decreases due to increasing effective nuclear charge. Lithium (Li) is in the second period and first group, while aluminum (Al) is in the third period and the thirteenth group. Aluminum has more electron shells in comparison to lithium, but it also has a higher effective nuclear charge due to the increased number of protons. So, the Al atom is smaller than the Li atom. This makes statement e true. In conclusion, statements a and e are true, while statements b, c, and d are false.

Key concepts

Cations vs Parent Atoms

When an atom forms a cation, it loses one or more electrons. As electrons are negatively charged, their removal results in a decrease in the overall electron-electron repulsion within the atom. This decrease causes the remaining electrons to be drawn closer to the nucleus. As a result, cations are generally smaller than their parent atoms because the effective nuclear charge (the pull from the nucleus on the electrons) is now concentrated over fewer electrons. Therefore, the statement "Cations are larger than their parent atoms" is false. Here's why:

• Loss of electrons results in decreased electron-electron repulsion.
• Remaining electrons are pulled closer to the nucleus.
• Smaller size due to fewer electrons surrounding the nucleus.

Electron Configuration

Electron configuration is a way to describe the arrangement of electrons in an atom. It is organized by the principle energy levels and subshells that electrons occupy. Understanding electron configuration helps explain why energy is required to remove an electron. Taking the example of a lithium atom, which has an electron configuration of 1s² 2s¹:

• The 1s electrons are closer to the nucleus and shield part of the nuclear charge from the outer 2s electron.
• Due to this shielding effect, less energy is required to remove the 2s electron compared to if it was not shielded.

However, removing an electron from an atom is an endothermic process, meaning energy must be supplied to overcome the attraction between the electron and the nucleus. Hence, the removal of an electron from lithium requires energy input, contradicting the notion that it is an exothermic process.

Periodic Table Trends

Periodic table trends play a crucial role in predicting the behavior of elements, including ionization energy. As you move across a period from left to right:

• Ionization energy increases because the effective nuclear charge is stronger; thus, electrons are held more tightly.
• Atomic size typically decreases due to stronger nuclear attraction pulling electrons closer.

For example, fluorine (F), located to the right of lithium (Li) in the same period, has a larger first ionization energy because the increased effective nuclear charge holds its electrons more tightly. In contrast, moving down a group, the ionization energy decreases because additional electron shells decrease the effective nuclear charge felt by the outermost electrons, and the atoms become larger, making it easier to remove an outer electron.

Atomic and Ionic Sizes

The size of atoms and ions can be influenced by several factors, such as the number of electron shells and effective nuclear charge. Atomic size decreases across a period due to the increase in nuclear charge, which pulls electrons closer to the nucleus.

• For instance, despite elements like lithium (Li) and aluminum (Al) being in different periods, aluminum's higher nuclear charge reduces its size compared to lithium, even though it has more electron shells.

In terms of ionic size, the loss of all electrons can lead to drastic changes. A hydrogen ion (H⁺) with no electrons contrasts starkly with a helium atom (He) which has two electrons, making the ion significantly smaller than the atom.