Question

(a) What is the oxidation state of \(\mathrm{P}\) in \(\mathrm{PO}_{4}{ }^{3-}\) and of \(\mathrm{N}\) in \(\mathrm{NO}_{3}^{-} ?(\mathbf{b})\) Why doesn't \(\mathrm{N}\) form a stable \(\mathrm{NO}_{4}^{3-}\) ion analogous to \(\mathrm{P}\) ?

Short answer

(a) The oxidation state of P in \(PO_{4}^{3-}\) is +5, and the oxidation state of N in \(NO_{3}^{-}\) is +5. (b) Nitrogen doesn't form a stable \(NO_{4}^{3-}\) ion due to its smaller atomic size, its tendency to form triple and double bonds, and lower electronegativity compared to phosphorus.

Step-by-step solution

Determine the oxidation state of P in PO₄³⁻

To determine the oxidation state of P in PO₄³⁻, we use the fact that the sum of the oxidation states of all atoms in a compound equals its overall charge. We know that the oxidation state of oxygen is typically -2. Since there's 4 oxygen atoms, the sum of their oxidation states is -8. Let x be the oxidation state of P. Therefore: \[x + (-8) = -3\]

Solve for x

Now, we can simply solve for x: \[x - 8 = -3\] \[x = -3 + 8\] \[x = 5\] The oxidation state of P in PO₄³⁻ is +5.

Determine the oxidation state of N in NO₃⁻

Similarly, to determine the oxidation state of N in NO₃⁻, we know that the sum of the oxidation states of all atoms in a compound equals its overall charge. We use the fact that the oxidation state of oxygen is typically -2. Since there's 3 oxygen atoms, the sum of their oxidation states is -6. Let y be the oxidation state of N. Therefore: \[y + (-6) = -1\]

Solve for y

Now, we can simply solve for y: \[y - 6 = -1\] \[y = -1 + 6\] \[y = 5\] The oxidation state of N in NO₃⁻ is +5.

Explain why N doesn't form a stable NO₄³⁻ ion

The fact that nitrogen does not form a stable NO₄³⁻ ion is due to a few factors. One reason is that nitrogen has a smaller atomic size compared to phosphorus, having only 2 electron shells to phosphorus's 3. This makes forming the NO₄³⁻ ion extremely difficult, as nitrogen would need to share more bonding electrons, which is unfavorable due to nitrogen's smaller size and its electron-electron repulsion. Another reason is that nitrogen has a tendency to form triple and double bonds, resulting in the more stable NO₂ and N₂O compounds. Lastly, nitrogen is less electronegative than phosphorus, making the bonding of NO₄³⁻ weaker, ultimately contributing to its unstable nature. In summary: (a) The oxidation state of P in PO₄³⁻ is +5, and the oxidation state of N in NO₃⁻ is +5. (b) Nitrogen doesn't form a stable NO₄³⁻ ion due to its smaller atomic size, its tendency to form triple and double bonds, and lower electronegativity compared to phosphorus.

Key concepts

Redox Reactions

Redox reactions are fundamental processes in chemistry where the oxidation states of atoms change. These reactions involve the transfer of electrons between species, underpinning many important chemical phenomena such as combustion, corrosion, and even respiration. To understand redox reactions, one must grasp the concept of oxidation states, which indicates the degree of oxidation of an atom in a molecule.

For instance, in the exercise, we determined the oxidation states of phosphorus (\textbf{P}) in \textbf{PO}\(_4^{3-}\) and nitrogen (\textbf{N}) in \textbf{NO}\(_3^{-}\) to be +5. By knowing the oxidation states, we can deduce that a reduction or oxidation has occurred if these states change in a chemical reaction. To identify redox reactions, look for changes in oxidation states and the movement of electrons from one atom or molecule to another.

Key signs of redox reactions include:

• Elements that change their oxidation state during the reaction.
• Formation of new compounds with different electron configurations.
• Transfer of electrons, which can sometimes be seen in a discharge or other electrical activity.

Grasping the concept of oxidation states not only helps in identifying redox reactions but also in balancing chemical equations for such reactions, which is a fundamental skill in chemistry.

Chemical Bonding

Chemical bonding is the force that holds atoms together in compounds. It's a complex interaction involving the attraction between electrons of one atom and the nucleus of another. In the educational exercise given, understanding chemical bonding helps explain why \textbf{N} does not form a stable \textbf{NO}\(_4^{3-}\) ion. There are several types of chemical bonds, with the main ones being ionic, covalent, and metallic bonds.

The stability of a chemical bond often depends on several factors, including atomic size, electron configuration, and electronegativity. For example, nitrogen's reluctance to form stable \textbf{NO}\(_4^{3-}\) ions can be attributed to its small atomic size and strong tendency to form double or triple bonds, which are typical characteristics of covalent bonding. Electronegativity also plays a crucial role, with less electronegative atoms being less likely to share electrons, leading to weaker bonds and less stable compounds.

Characteristics of Chemical Bonds

• Ionic Bonds: Formed when an atom donates an electron to another atom, typically between metals and non-metals.
• Covalent Bonds: When atoms share electrons equally or unequally, primarily between non-metal atoms.
• Metallic Bonds: The delocalized 'sea of electrons' that move freely around metal ions, characteristic of metals.

These bonding patterns are key to predicting the structure, reactivity, and properties of molecules.

Periodic Table Properties

The periodic table is a comprehensive layout of elements organized by increasing atomic number, electron configuration, and recurring chemical properties. Elements are grouped into periods and groups, which reflect their properties and trends. By studying the periodic table properties, we can make educated guesses regarding the behavior of elements in chemical reactions, like oxidation and bonding preferences.

Two properties highlighted in the exercise are atomic size and electronegativity. Atomic size decreases across a period and increases down a group. Nitrogen, being in the second period, has a smaller atomic size than phosphorus, which is in the third period. This impacts its ability to form stable compounds. Electronegativity, which is the ability of an atom to attract and hold onto electrons, varies across the periodic table as well.

Periodic Trends

• Electronegativity: Generally increases from left to right across a period and decreases down a group.
• Atomic Size: Decreases from left to right across a period due to increasing nuclear charge, and increases down a group as new electron shells are added.
• Ionization Energy: The energy required to remove an electron, increases across a period and decreases down a group.

These trends are essential for understanding why certain elements behave differently in chemical reactions and help in predicting the formation of stable or unstable ions, such as \textbf{NO}\(_4^{3-}\) as compared to \textbf{PO}\(_4^{3-}\). Learning how to navigate the periodic table and use it as a tool is an invaluable skill in the study of chemistry.