Question

Write the electron configurations for the following ions, and determine which have noble-gas configurations: \((\mathbf{a})\mathrm{Ru}^{3+}\) \((\mathbf{b}) \mathrm{As}^{3-},(\mathbf{c}) \mathrm{Y}^{3+},(\mathbf{d}) \mathrm{Pd}^{2+},(\mathbf{e}) \mathrm{Pb}^{2+},(\mathbf{f}) \mathrm{Au}^{3+}.\)

Short answer

The electron configurations for the given ions are as follows: (a) \(Ru^{3+}\): \(1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6 \ 4d^{7}\) - not a noble gas configuration (b) \(As^{3-}\): \(1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6\) - matches Kr (noble gas configuration) (c) \(Y^{3+}\): \(1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6\) - matches Kr (noble gas configuration) (d) \(Pd^{2+}\): \(1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6 \ 4d^{8}\) - not a noble gas configuration (e) \(Pb^{2+}\): \(1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6 \ 4d^{10} \ 4f^{14} \ 5s^2 \ 5p^6 \ 5d^{10} \ 6s^2 \ 6p^4\) - not a noble gas configuration (f) \(Au^{3+}\): \(1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6 \ 4d^{10} \ 4f^{14} \ 5s^2 \ 5p^6 \ 5d^{9} \ 6s^1\) - not a noble gas configuration Out of these ions, only \(As^{3-}\) and \(Y^{3+}\) have noble-gas configurations (matching Kr).

Step-by-step solution

Identify the atomic number of each ion

Look at the periodic table and find the atomic number (number of protons) of the neutral atom. This will also give us the number of electrons in the neutral atom since the number of protons equals the number of electrons in a neutral atom.

Determine the number of electrons in the ion

Calculate the number of electrons in the ion by adding or subtracting the given charge from the atomic number of the neutral atom.

Write the electron configuration for each ion

Based on the number of electrons, write down the electron configuration for each ion, following the filling order of subshells.

Check if the electron configurations match any noble gases

Look at the periodic table and see if any of the ions' electron configurations match those of noble gases. Now, let's go through each ion and write their electron configurations.

Electron configuration of Ru^{3+}

Ru (Ruthenium) has an atomic number of 44. Ru^{3+} ion means three electrons are removed from the neutral atom, so we have 44 - 3 = 41 electrons. The electron configuration will be: \[1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6 \ 4d^{7}\] This does not match any noble gas configuration.

Electron configuration of As^{3-}

As (Arsenic) has an atomic number of 33. As^{3-} ion means three electrons are added to the neutral atom, so we have 33 + 3 = 36 electrons. The electron configuration will be: \[1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6\] This matches the electron configuration of a noble gas: Kr (Krypton).

Electron configuration of Y^{3+}

Y (Yttrium) has an atomic number of 39. Y^{3+} ion means three electrons are removed from the neutral atom, so we have 39 - 3 = 36 electrons. The electron configuration will be: \[1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6\] This matches the electron configuration of a noble gas: Kr (Krypton).

Electron configuration of Pd^{2+}

Pd (Palladium) has an atomic number of 46. Pd^{2+} ion means two electrons are removed from the neutral atom, so we have 46 - 2 = 44 electrons. The electron configuration will be: \[1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6 \ 4d^{8}\] This does not match any noble gas configuration.

Electron configuration of Pb^{2+}

Pb (Lead) has an atomic number of 82. Pb^{2+} ion means two electrons are removed from the neutral atom, so we have 82 - 2 = 80 electrons. The electron configuration will be: \[1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6 \ 4d^{10} \ 4f^{14} \ 5s^2 \ 5p^6 \ 5d^{10} \ 6s^2 \ 6p^4\] This does not match any noble gas configuration.

Electron configuration of Au^{3+}

Au (Gold) has an atomic number of 79. Au^{3+} ion means three electrons are removed from the neutral atom, so we have 79 - 3 = 76 electrons. The electron configuration will be: \[1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ 3d^{10} \ 4s^2 \ 4p^6 \ 4d^{10} \ 4f^{14} \ 5s^2 \ 5p^6 \ 5d^{9} \ 6s^1\] This does not match any noble gas configuration. From the electron configurations, we can see that only the As^{3-} and Y^{3+} ions have noble-gas configurations (matching Kr).

Key concepts

Noble Gas Configuration

Noble gas configuration refers to an atom or ion whose electron configuration is the same as one of the noble gases, located in the far right column of the periodic table. These gases include helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn).
Noble gases are known for their stability because they have a full outer electron shell, which is the driving factor for most elements when they gain or lose electrons.
When ions such as \( ext{As}^{3-}\) and \( ext{Y}^{3+}\) achieve the electron configuration of krypton, they have reached a state of maximum stability similar to that of krypton. This is why these configurations are often sought after when atoms form ions.

This configuration provides an insight into why certain elements tend to gain or lose specific numbers of electrons to achieve this stability. By comparing to noble gas electron configurations, we can determine the most stable form for many ions.

Periodic Table

The periodic table is a crucial tool in chemistry that organizes all known elements according to increasing atomic numbers and similar chemical properties.
The elements are arranged in rows called periods and columns known as groups or families.
The location of elements on the periodic table can tell us about the element's electron configuration and how it will interact with other elements.

The periodic table helps us understand which electrons are involved in chemical reactions by displaying elements in such a way that reflects recurring trends called periodicity. For instance, elements in the same group typically possess similar chemical properties and electron configurations in their outermost shells, making them behave similarly in chemical situations.

• The first column contains alkali metals, which have one electron in their outermost shell and are highly reactive.
• The last column consists of noble gases, which have full outer shells making them highly stable and generally non-reactive.
• The transition metals, found in the middle, have partially filled d subshells and exhibit unique properties like variable oxidation states.

The periodic table is not just a way to organize elements; it's a map to predict how they behave.

Ionization

Ionization is the process by which an atom or molecule acquires a positive or negative charge by gaining or losing electrons to form ions.
It involves the energy needed to remove one or more electrons from a neutral atom, transforming it into an ion.
Ionization is crucial in understanding how elements form compounds and how they can exist in charged forms, like \( ext{Ru}^{3+}\) or \( ext{As}^{3-}\).

Different elements require varying amounts of energy to remove electrons based on their electron configurations. This concept is known as ionization energy, which influences the reactivity and chemical behavior of elements.

• Elements with a full outer electron shell, such as the noble gases, have high ionization energies since they are already stable.
• Elements in groups 1 and 2 generally have lower ionization energies, making them more likely to lose electrons and form positive ions.

Gaining an understanding of ionization helps predict which electrons an element is likely to lose or gain during chemical reactions.

Subshells Filling Order

The subshells filling order is an essential concept in understanding the electron configuration of atoms. It explains the sequence in which electrons populate the available subshells in an atom.
Electrons fill subshells following a specific order that minimizes the energy of the atom, dictated by the Aufbau principle, Hund's rule, and the Pauli exclusion principle.

Following the order of filling, the subshells are populated in increasing energy levels, with the s subshell filled before p, which is filled before d, followed by f as described by the sequence: \(1s^2\, 2s^2\, 2p^6\, 3s^2\, 3p^6\, 4s^2\, 3d^{10}\, 4p^6, \)... and so on.
The correct filling order helps predict how electrons are arranged within an atom or ion:

• Electrons fill the 1s orbital first since it has the lowest energy.
• They then proceed to 2s, followed by 2p, and continue according to the sequence as more electrons are added.
• This order ensures atoms settle in their lowest possible energy states, which is key to understanding their chemical properties.

This sequence explains why the configurations of certain ions, such as in the exercise, may or may not resemble a noble gas configuration. Understanding the subshell filling order is crucial for constructing accurate electron configurations.