Question

What is the electron configuration for the transition metal ion(s) in each of the following compounds? a. \(\left(\mathrm{NH}_{4}\right)_{2}\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2} \mathrm{Cl}_{4}\right]\) b. \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{2}\left(\mathrm{NH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}\right)_{2}\right] \mathrm{I}_{2}\) c. \(\mathrm{Na}_{2}\left[\mathrm{TaF}_{7}\right]\) d. \(\left[\mathrm{Pt}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{I}_{2}\right]\left[\mathrm{PtI}_{4}\right]\) Pt forms \(+2\) and \(+4\) oxidation states in compounds.

Short answer

The electron configurations of the transition metal ions in the given compounds are: a. Fe (+2 oxidation state): \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^6\] b. Co (+3 oxidation state): \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^6\] c. Ta (+5 oxidation state): \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^0 6s^2\] d. Pt (+2 oxidation state): \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^8 6s^1\] Pt (+4 oxidation state): \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^7\]

Step-by-step solution

Identify the transition metal ion and its oxidation state

For each compound, we need to identify the transition metal ion and also the oxidation state of the transition metal ion. a. The transition metal ion in the first compound is Fe. b. The transition metal ion in the second compound is Co. c. The transition metal ion in the third compound is Ta. d. The transition metal ion in the fourth compound is Pt (appearing twice).

Determine the oxidation state of the transition metal ion

a. In \(\left(\mathrm{NH}_{4}\right)_{2}\left[\mathrm{Fe}\left(\mathrm{H}_{2}\mathrm{O}\right)_{2} \mathrm{Cl}_{4}\right]\), the oxidation state of Fe is +2. b. In \(\left[\mathrm{Co}\left(\mathrm{NH}_{3}\right)_{2}\left(\mathrm{NH}_{2}\mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}\right)_{2}\right] \mathrm{I}_{2}\), the oxidation state of Co is +3. c. In \(\mathrm{Na}_{2}\left[\mathrm{TaF}_{7}\right]\), the oxidation state of Ta is +5. d. In \(\left[\mathrm{Pt}\left(\mathrm{NH}_{3}\right)_{4}\mathrm{I}_{2}\right]\left[\mathrm{PtI}_{4}\right]\): - For the \(\left[\mathrm{Pt}\left(\mathrm{NH}_{3}\right)_{4}\mathrm{I}_{2}\right]\), the oxidation state of Pt is +2. - For the \(\left[\mathrm{PtI}_{4}\right]\), the oxidation state of Pt is +4.

Determine the electron configuration of transition metal ions

Now that we have the oxidation state of the transition metal ions, we can find their electron configurations using the periodic table. a. Fe has an atomic number of 26 and an electron configuration of: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^6 4s^2\] In the +2 oxidation state, Fe loses 2 electrons, giving the electron configuration: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^6\] b. Co has an atomic number of 27 and an electron configuration of: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^7 4s^2\] In the +3 oxidation state, Co loses 3 electrons, giving the electron configuration: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^6\] c. Ta has an atomic number of 73, and an electron configuration of: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^3 6s^2\] In the +5 oxidation state, Ta loses 5 electrons, giving the electron configuration: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^0 6s^2\] d. Pt has an atomic number of 78, and an electron configuration of: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^9 6s^1\] In the +2 oxidation state, Pt loses 2 electrons as: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^8 6s^1\] In the +4 oxidation state, Pt loses 4 electrons as: \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^7\]

Key concepts

Oxidation states

Oxidation states, also known as oxidation numbers, reflect the degree of oxidation or reduction of an element in a chemical compound. They are essential for understanding how electrons are distributed in a molecule or ion. The oxidation state is particularly significant when it comes to transition metals, as it can affect the compound's physical and chemical properties.

In our example exercise, the transition metals like Iron (Fe), Cobalt (Co), Tantalum (Ta), and Platinum (Pt) sport various oxidation states. For instance, Fe can exhibit an oxidation state of +2, which means it loses two electrons from its atomic configuration. Similarly, Pt can have oxidation states of +2 or +4 in different compounds.

Understanding these states is crucial as they dictate aspects like the color of the compound, its magnetic properties, and its reactivity with other chemicals. Oxidation states are determined based on specific rules, such as the fact that the sum of oxidation states in a neutral compound must equal zero.

Transition metals

Transition metals occupy the d-block in the periodic table and are known for having variable oxidation states. This is because they can lose varying numbers of their d and sometimes s-electrons. This feature of transition metals gives rise to rich chemistry, inclusive of colorful compounds and catalytic behaviors.

In the given exercise, transition metals like Fe, Co, Ta, and Pt form the backbone of the complexes. Due to their ability to change oxidation state, these metals can form various ionic compounds that have different properties. Transition metals also typically display metallic properties such as high conductivity, malleability, and a bright luster.

Moreover, the ability of these metals to form stable complex ions with varying numbers of ligands — molecules that donate pairs of electrons to the metal — further demonstrates their versatile chemistry. Properties such as magnetism and an expansive range of colors also arise from the differing oxidation states and electron configurations.

Periodic table

The periodic table is an organized chart that arranges elements by increasing atomic number and groups them by shared properties. It's divided into various blocks, such as the s-block, p-block, d-block, and f-block, which relate to the electron configuration of the elements they house.

Transition metals are found in the d-block of the periodic table, typically spanning from groups 3 to 12. The position of elements in the periodic table gives insight into their electron configurations, which in turn informs their chemical behavior. For instance, familiarity with the table allows one to deduce that elements like Fe, Co, and Pt are transition metals with characteristics influenced by partially filled d-orbitals.

Each row of the periodic table — or period — corresponds to the filling of a principal energy level of electrons, which helps dictate the pattern of chemical reactivity and properties. The periodic nature of the table allows predictions of elemental properties and reactivity, which is incredibly useful in chemical analyses and synthesis.

Electron configuration of ions

The electron configuration of ions refers to the arrangement of electrons in an atom's orbitals after it has lost or gained electrons to form an ion. This process involves either the removal or addition of electrons to achieve a more stable electron arrangement.

For the transition metals examined in the exercise, determining the electron configuration of their ions begins by establishing the number of electrons in their neutral state. From there, electrons are then removed in accordance with their oxidation states. For example, Fe in its +2 oxidation state has two fewer electrons than the neutral Fe atom, leading to a partially filled 3d subshell.

This change in configuration is significant as it influences the chemical behavior of the ion. Electrons are removed from the outermost shell, typically starting from the 4s orbital before the 3d in transition metal ions. Understanding this sequence of removal aids in predicting ionic charges and bond formation in various chemical reactions.