Chapter 3: Problem 29
The number of unpaired electrons present in \(\mathrm{Fe}^{+3}\) is
Short Answer
Expert verified
Answer: The Fe+3 ion contains 5 unpaired electrons.
Step by step solution
01
Determine the atomic number of Fe
The atomic number of Iron (Fe) can be found on the periodic table. It is 26.
02
Write the electron configuration of the Fe atom
Since the atomic number of Fe is 26, it has 26 electrons. The electron configuration for Fe can be determined by following the order of orbitals: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. The electron configuration of Fe is: \(\mathrm{1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 4s^2 \ 3d^6}\)
03
Determine the electron configuration of the \(\mathrm{Fe^{+3}}\) ion
Since \(\mathrm{Fe^{+3}}\) has lost three electrons, we need to remove three electrons from the electron configuration. Always remove electrons from the highest energy level first (4s, then 3d). The electron configuration of \(\mathrm{Fe^{+3}}\) is: \(\mathrm{1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 4s^0 \ 3d^5}\)
04
Count the number of unpaired electrons
The electrons in the 1s, 2s, 2p, 3s, and 3p orbitals are all paired. The only unpaired electrons are in the 3d orbital. There are 5 unpaired electrons in the \(\mathrm{Fe^{+3}}\) ion.
The number of unpaired electrons present in \(\mathrm{Fe^{+3}}\) is 5.
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Unpaired Electrons
Unpaired electrons are those that exist alone in an atomic orbital. Electrons naturally pair up because paired electrons have opposite spins, which stabilizes the atom. However, in some cases, electrons do not find a pair and remain unpaired.
Understanding the significance of unpaired electrons is crucial because they are often implicated in the magnetic properties of an atom. Atoms with unpaired electrons are typically paramagnetic, meaning they are attracted to magnetic fields, while those without unpaired electrons are considered diamagnetic. In the exercise, we see that the Fe^{+3} ion has five unpaired electrons, which can significantly influence its behavior in magnetic fields.
In chemistry, determining the number of unpaired electrons is essential for predicting the reactivity and bonding of an element. For example, the unpaired electrons in the iron ion allow it to participate in the formation of chemical bonds, particularly in complex compounds such as those found in biological systems.
Understanding the significance of unpaired electrons is crucial because they are often implicated in the magnetic properties of an atom. Atoms with unpaired electrons are typically paramagnetic, meaning they are attracted to magnetic fields, while those without unpaired electrons are considered diamagnetic. In the exercise, we see that the Fe^{+3} ion has five unpaired electrons, which can significantly influence its behavior in magnetic fields.
In chemistry, determining the number of unpaired electrons is essential for predicting the reactivity and bonding of an element. For example, the unpaired electrons in the iron ion allow it to participate in the formation of chemical bonds, particularly in complex compounds such as those found in biological systems.
Periodic Table
The periodic table is an organized chart of elements arranged in order of increasing atomic number. The atomic number is synonymous with the number of protons in an atom's nucleus. The table allows chemists to understand trends and predict the properties of elements.
Elements in the same column, or group, typically have similar chemical behaviors due to their similar valence electron configurations. For example, within the exercise context, iron (Fe) is in the transition metals group, characterized by their d-orbitals being filled with electrons.
Utilizing the periodic table, we can locate iron and immediately determine its atomic number, which is crucial when writing electron configurations or understanding how an atom will react chemically. The table is also invaluable in studying and organizing elements, as it shows relationships and patterns among them.
Elements in the same column, or group, typically have similar chemical behaviors due to their similar valence electron configurations. For example, within the exercise context, iron (Fe) is in the transition metals group, characterized by their d-orbitals being filled with electrons.
Utilizing the periodic table, we can locate iron and immediately determine its atomic number, which is crucial when writing electron configurations or understanding how an atom will react chemically. The table is also invaluable in studying and organizing elements, as it shows relationships and patterns among them.
Atomic Number
The atomic number is a fundamental characteristic of an element. It represents the number of protons found in the nucleus of an atom and by extension, also indicates the number of electrons in a neutral atom. Each element has a unique atomic number, making it an identifying feature.
In our exercise, the atomic number for iron (Fe) is 26. This is directly taken from the periodic table. The atomic number not only tells us how many protons are in the nucleus but also guides us to determine the number of electrons and their standard arrangement around the nucleus before any ionization occurs, as in the case of Fe^{+3}. Understanding atomic number facilitates predictions about an element’s chemical behavior and its position on the periodic table.
In our exercise, the atomic number for iron (Fe) is 26. This is directly taken from the periodic table. The atomic number not only tells us how many protons are in the nucleus but also guides us to determine the number of electrons and their standard arrangement around the nucleus before any ionization occurs, as in the case of Fe^{+3}. Understanding atomic number facilitates predictions about an element’s chemical behavior and its position on the periodic table.
Orbital Theory
Orbital theory is a fundamental concept in chemistry that describes the regions around an atom where electrons are most likely to be found. These regions are known as orbitals, and they come in different shapes and sizes, such as s, p, d, and f. Each orbital can hold a maximum of two electrons.
In our iron example, we consider the 3d orbital where the unpaired electrons reside. According to the Aufbau principle, electrons fill the lowest energy orbitals first; however, once we reach d and f orbitals, things become more complex due to their higher energy levels.
The distribution of electrons among an atom's orbitals determines its electron configuration, which in turn affects the atom's chemical properties. In ions such as Fe^{+3}, their orbital occupancy changes due to the loss of electrons, which then defines the number of unpaired electrons, showing a deep connection between orbital theory and other concepts like atomic number and the periodic table.
In our iron example, we consider the 3d orbital where the unpaired electrons reside. According to the Aufbau principle, electrons fill the lowest energy orbitals first; however, once we reach d and f orbitals, things become more complex due to their higher energy levels.
The distribution of electrons among an atom's orbitals determines its electron configuration, which in turn affects the atom's chemical properties. In ions such as Fe^{+3}, their orbital occupancy changes due to the loss of electrons, which then defines the number of unpaired electrons, showing a deep connection between orbital theory and other concepts like atomic number and the periodic table.
