Chapter 3: Problem 53
Which of the following elements will have highest second ionisation enthalpy? (a) \(15^{2} 2 s^{2} 2 p^{6} 3 s^{2}\) (b) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{1}\) (c) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{2}\) (d) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{3}\)
Short Answer
Expert verified
The element with the configuration (a) will have the highest second ionisation enthalpy.
Step by step solution
01
Understanding Second Ionisation Enthalpy
Second ionisation enthalpy refers to the amount of energy required to remove a second electron from a singly charged cation. It typically increases across a period as the atomic radius decreases and nuclear charge increases, making it harder to remove an electron. It also increases from a filled or half-filled subshell as compared to one which has just started to fill up because filled and half-filled subshells are more stable.
02
Analyze Each Configuration
We need to determine which configuration will make it the hardest to remove another electron. In general, atoms with half-filled or completely filled subshells will have higher second ionisation enthalpy due to their more stable electron configurations. Analyze the given electron configurations to see which represents an atom with either a filled or half-filled subshell after the removal of the first electron.
03
Comparing Electron Configurations
(a) After removing one electron, we would have a noble gas configuration, which is already stable. (b) Removing one electron from this atom would leave a single electron in the 3s orbital, which would not be very stable. (c) After removing one electron, the atom would have a filled 3s subshell but would not be at a half-filled stability point. (d) After removing one electron, this atom would have 3s2 3p2 configuration, which isn't at a half-filled or filled subshell. Therefore, none of these configurations would lead to half-filled or fully filled subshell stability after removing one electron.
04
Identify the Highest Second Ionisation Enthalpy
Among the options given, after the removal of one electron, an atom with a noble gas configuration would be the most stable and thus require the most energy to remove the next electron. This corresponds to option (a). Thus, the element represented by (a) will have the highest second ionization enthalpy because after losing one electron, it attains a noble gas configuration, which is extremely stable.
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Ionisation Energy
Ionisation energy is the amount of energy required to remove an electron from an atom in its gaseous state. The term second ionisation energy specifically refers to the energy needed to remove the second electron after the first has already been removed, thereby creating a +2 ion. Understanding ionisation energy is crucial because it provides insight into an atom's reactivity and chemical behavior.
Key factors affecting ionisation energy include atomic size, nuclear charge, and the electron shielding effect. Generally, the smaller the atom and the greater the nuclear charge, the higher the ionisation energy. Electrons closer to the nucleus experience a stronger attraction and therefore require more energy to be removed. As you move across a period in the periodic table, the ionisation energy tends to increase because atoms have more protons (increasing nuclear charge) but similar shielding effect. This is because the added electrons occupy the same energy shell, which does not substantially increase the distance from the nucleus.
Key factors affecting ionisation energy include atomic size, nuclear charge, and the electron shielding effect. Generally, the smaller the atom and the greater the nuclear charge, the higher the ionisation energy. Electrons closer to the nucleus experience a stronger attraction and therefore require more energy to be removed. As you move across a period in the periodic table, the ionisation energy tends to increase because atoms have more protons (increasing nuclear charge) but similar shielding effect. This is because the added electrons occupy the same energy shell, which does not substantially increase the distance from the nucleus.
Electron Configuration
Electron configuration describes how electrons are distributed in an atom's orbitals. Electrons are arranged in shells around the nucleus, and each shell can hold a specific number of electrons. Within shells, there are subshells named s, p, d, and f, each with a different capacity for electrons.
Understanding the rules of electron configuration is fundamental. The Aufbau principle dictates that electrons fill orbitals starting with the lowest energy level before moving to higher ones. Hund's rule states that electrons will fill an unoccupied orbital before they pair up. Finally, Pauli's exclusion principle maintains that no two electrons in the same atom can have identical quantum numbers, ensuring that electron pairs within the same orbital have opposite spins. These principles result in patterns of electron configuration that underpin periodic trends and the chemical properties of elements.
Understanding the rules of electron configuration is fundamental. The Aufbau principle dictates that electrons fill orbitals starting with the lowest energy level before moving to higher ones. Hund's rule states that electrons will fill an unoccupied orbital before they pair up. Finally, Pauli's exclusion principle maintains that no two electrons in the same atom can have identical quantum numbers, ensuring that electron pairs within the same orbital have opposite spins. These principles result in patterns of electron configuration that underpin periodic trends and the chemical properties of elements.
Atomic Structure
Atomic structure refers to the arrangement of particles within an atom, which includes protons, neutrons, and electrons. Protons and neutrons form the atom's nucleus, and electrons orbit this nucleus, largely defining the atom's properties and behavior. The atomic number is the number of protons in the nucleus and determines the element's identity, while the number of electrons affects an atom's chemical properties.
As you delve into atomic structure, you'll understand that the arrangement of electrons is not random. They are organized into shells and subshells based on their energy levels. This is why when you remove electrons, you typically take them from the outermost shell, as these electrons are less tightly bound to the nucleus. The concept of atomic structure is pivotal in explaining why certain elements have high second ionisation energies, since removing an electron from a stable, filled or half-filled shell requires more energy.
As you delve into atomic structure, you'll understand that the arrangement of electrons is not random. They are organized into shells and subshells based on their energy levels. This is why when you remove electrons, you typically take them from the outermost shell, as these electrons are less tightly bound to the nucleus. The concept of atomic structure is pivotal in explaining why certain elements have high second ionisation energies, since removing an electron from a stable, filled or half-filled shell requires more energy.
Periodic Trends
Periodic trends refer to patterns in the periodic table that can predict various element properties, including ionisation energy, atomic radius, electronegativity, and metallic character. These trends are the result of electron configuration and the effective nuclear charge experienced by the outermost electrons.
When discussing ionisation energies, the trend generally increases across a period, as nuclear charge increases and electrons are added to the same energy level without significantly increasing electron shielding. Conversely, ionisation energy tends to decrease down a group as the atomic number increases, because additional shells of electrons are added, which increases the distance from the nucleus and electron shielding, both of which weaken the attraction between the nucleus and the outer electrons.
Understanding periodic trends is essential for predicting chemical reactivity and the amounts of energy involved in ionisation processes. For instance, within a period, elements on the left have lower ionisation energies than those on the right, reflecting the increased nuclear charge and decreased atomic radius as one moves to the right.
When discussing ionisation energies, the trend generally increases across a period, as nuclear charge increases and electrons are added to the same energy level without significantly increasing electron shielding. Conversely, ionisation energy tends to decrease down a group as the atomic number increases, because additional shells of electrons are added, which increases the distance from the nucleus and electron shielding, both of which weaken the attraction between the nucleus and the outer electrons.
Understanding periodic trends is essential for predicting chemical reactivity and the amounts of energy involved in ionisation processes. For instance, within a period, elements on the left have lower ionisation energies than those on the right, reflecting the increased nuclear charge and decreased atomic radius as one moves to the right.