Question

Identify each statement as true or false: (a) lonization energies are always endothermic. (b) Potassium has a larger first ionization energy than lithium. (c) The second ionization energy of the sodium atom is larger than the second ionization energy of the magnesium atom. (d) The third ionization energy is three times the first ionization energy of an atom.

Short answer

(a) True: Ionization energies are always endothermic since they require energy input to remove an electron from a gaseous atom or ion. (b) False: Potassium has a lower first ionization energy than lithium because it is located below lithium in the same group in the periodic table. (c) False: The second ionization energy of the sodium atom is smaller than the second ionization energy of the magnesium atom due to the higher effective nuclear charge in magnesium. (d) False: There is no fixed relationship between the third ionization energy and the first ionization energy of an element. The difference is specific to each element and its electron configuration.

Step-by-step solution

Statement a: Ionization energies are always endothermic.

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The process of ionization requires energy input to overcome the attractive forces between the negatively charged electron and the positively charged nucleus. Since energy is being absorbed to remove the electron, the process is endothermic. Thus, the statement is true.

Statement b: Potassium has a larger first ionization energy than lithium.

When moving across a period in the periodic table, ionization energy generally increases due to a greater effective nuclear charge. Lithium is in the second period, while potassium is in the fourth period. When moving down a group, ionization energy typically decreases because of an increase in atomic size and shielding effect. Since potassium is located below lithium in the same group, its first ionization energy is lower than that of lithium. Therefore, the statement is false.

Statement c: The second ionization energy of the sodium atom is larger than the second ionization energy of the magnesium atom.

Sodium and magnesium are both in the third period, with magnesium being to the right of sodium. The second ionization energy of an element corresponds to the energy required to remove the second electron. In the case of sodium, the second electron is being removed from a complete \(2p^6\) subshell, while in magnesium, it is being removed from an almost filled \(3s^2\) subshell. The effective nuclear charge of magnesium is greater than that of sodium, making it more difficult to remove the second electron. Therefore, the second ionization energy of magnesium is larger than that of sodium, and the statement is false.

Statement d: The third ionization energy is three times the first ionization energy of an atom.

Ionization energy generally increases with the removal of each subsequent electron due to a smaller atomic size and a higher effective nuclear charge. However, there is no fixed relationship between the third ionization energy and the first ionization energy of an element. The difference is specific to each element and its electron configuration. As a result, the statement is false.

Key concepts

Periodic Table

The periodic table is an organized chart that displays all known chemical elements in an informative array. It is structured in order of increasing atomic number, meaning from left to right and top to bottom, the elements rise in the number of protons. This arrangement helps in determining various properties of elements, including ionization energy.
Each row of the periodic table is termed a period, while each column is known as a group. Elements within the same group often share similar chemical properties because they have the same number of valence electrons.

• Characteristics such as metallic character and atomic size change systematically as you move across or down the table.
• Ionization energy tends to increase across a period due to the greater effective nuclear charge.
• Conversely, it decreases down a group as the atomic size increases, making it easier to remove an electron.

Understanding this predictable pattern helps in predicting the behavior of elements, especially when it comes to reactions and bonding.

Electron Configuration

Electron configuration is a notation that shows the distribution of electrons across the orbitals of an atom. It is essential for understanding how elements will react chemically and what kinds of bonds they'll form.
Electrons fill orbitals in a specific sequence known as the Aufbau principle. This principle is guided by energy levels and sublevels, which dictate that electrons fill the lowest energy orbitals first.

• For example, hydrogen's electron configuration is simply \(1s^1\), meaning one electron in the lowest energy orbital.
• Magnesium, with a more complex configuration, is \(1s^2 2s^2 2p^6 3s^2\).

The electron configuration not only helps predict ionization energy but also gives insights on the reactivity and bonding patterns of the element. Recognizing configurations can explain why certain elements are more stable or more reactive than others.

Effective Nuclear Charge

Effective nuclear charge (") is a concept that explains the net positive charge experienced by electrons in the valence shell. It accounts for the shielding effect of inner electrons, which mask the full positive charge of the protons in the nucleus.
In an atom, electrons closer to the nucleus repel those further away, reducing the latter's exposure to the nuclear charge. This is known as electron shielding.

• As you move right across a period, the effective nuclear charge generally increases since the number of protons in the nucleus increases while the shielding effect stays relatively constant.
• This increase results in a stronger attraction between the nucleus and the outermost electrons, thus raising the ionization energy.

Effective nuclear charge is crucial for understanding trends in atomic size and reactivity across the periodic table.

Atomic Size

Atomic size, or atomic radius, refers to the distance from the nucleus of an atom to the outer boundary of its electron cloud. It is a key factor that affects how elements interact and bond with each other.
Atomic size increases as you move down a group due to the addition of electron shells. Each subsequent shell is further from the nucleus, outweighing the increase in nuclear charge. However, as you move across a period from left to right, the atomic size decreases:

• Increased effective nuclear charge pulls electrons closer to the nucleus.
• Therefore, despite having more electrons, the atoms become smaller because the added electrons enter the same shell.

Knowing about atomic size is useful for predicting trends in ionization energy, as smaller atoms generally have higher ionization energies, making it harder to remove electrons.

Endothermic Reaction

An endothermic reaction is a chemical reaction that absorbs energy from its surroundings in the form of heat. This concept is critical when discussing ionization energy because ionization is inherently an endothermic process.
When an atom loses an electron, it requires input energy to overcome the attraction between the negatively charged electron and the positively charged nucleus. Making this process endothermic means it does not occur spontaneously; instead, energy must be supplied.

• Ionization energy is always positive, reflecting this absorption of energy.
• Endothermic processes are crucial for understanding how reactions occur and for predicting the conditions needed to initiate them.

Comprehending the endothermic nature of ionization assists in predicting reaction behaviors and understanding energy changes during chemical processes.