Question

Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group \(6 \mathrm{~A}\) element in each: (a) sulfur tetrachloride, (b) selenium trioxide, (c) sodium thiosulfate, (d) hydrogen sulfide, (e) sulfuric acid, (f) sulfur dioxide, (g) mercury telluride.

Short answer

(a) Sulfur tetrachloride: \(\mathrm{SCl_4}\), oxidation state of S: +4 (b) Selenium trioxide: \(\mathrm{SeO_3}\), oxidation state of Se: +6 (c) Sodium thiosulfate: \(\mathrm{Na_2S_2O_3}\), oxidation state of central S: 0, oxidation state of oxygen-bound S: -2 (d) Hydrogen sulfide: \(\mathrm{H_2S}\), oxidation state of S: -2 (e) Sulfuric acid: \(\mathrm{H_2SO_4}\), oxidation state of S: +6 (f) Sulfur dioxide: \(\mathrm{SO_2}\), oxidation state of S: +4 (g) Mercury telluride: \(\mathrm{HgTe}\), oxidation state of Te: -2

Step-by-step solution

(a) Sulfur tetrachloride

Sulfur tetrachloride has one sulfur atom and four chlorine atoms. The chemical formula is therefore: \[ \mathrm{SCl_4} \] To find the oxidation state of sulfur, we can use the rule that the sum of the oxidation states of all the atoms in a neutral molecule must be zero. The oxidation state of chlorine is -1, and there are four chlorine atoms, so the total oxidation state of the chlorines is -4. Since the sum must be zero, the oxidation state of sulfur is: \[ \mathrm{Oxidation~state~of~S} = +4 \]

(b) Selenium trioxide

Selenium trioxide has one selenium atom and three oxygen atoms. The chemical formula is therefore: \[ \mathrm{SeO_3} \] To find the oxidation state of selenium, we can use the rule that the sum of the oxidation states of all the atoms in a neutral molecule must be zero. The oxidation state of oxygen is -2, and there are three oxygen atoms, so the total oxidation state of the oxygens is -6. Since the sum must be zero, the oxidation state of selenium is: \[ \mathrm{Oxidation~state~of~Se} = +6 \]

(c) Sodium thiosulfate

Sodium thiosulfate has two sodium atoms, one sulfur atom, and one additional sulfur atom bonded to three oxygen atoms. The chemical formula is therefore: \[ \mathrm{Na_2S_2O_3} \] We will now find the oxidation state for the sulfur atom. The oxidation state of sodium is +1 and the oxidation state of oxygen is -2. The sum of the oxidation states for the sodiums and the oxygens must be balanced by the two sulfur atoms. We can write this as follows: \[ +1(2) + (\mathrm{Oxidation~state~of~central~S}) + (\mathrm{Oxidation~state~of~oxygen-bound~S}) = 0 \] The central sulfur is assumed to start at an oxidation state of 0 since it is not bound to any electronegative elements. Therefore: \[ +1(2) + 0 + (\mathrm{Oxidation~state~of~oxygen-bound~S}) = 0 \] This yields an oxidation state of -2 for the oxygen-bound sulfur.

(d) Hydrogen sulfide

Hydrogen sulfide has two hydrogen atoms and one sulfur atom. The chemical formula is therefore: \[ \mathrm{H_2S} \] To find the oxidation state of sulfur, we can use the rule that the sum of the oxidation states of all the atoms in a neutral molecule must be zero. The oxidation state of hydrogen is +1, and there are two hydrogen atoms, so the total oxidation state of the hydrogens is +2. Since the sum must be zero, the oxidation state of sulfur is: \[ \mathrm{Oxidation~state~of~S} = -2 \]

(e) Sulfuric acid

Sulfuric acid has two hydrogen atoms, one sulfur atom, and four oxygen atoms. The chemical formula is therefore: \[ \mathrm{H_2SO_4}\] To find the oxidation state of sulfur, we can use the rule that the sum of the oxidation states of all the atoms in a neutral molecule must be zero. The oxidation state of hydrogen is +1 and the oxidation state of oxygen is -2. Since the sum must be zero, we can set up the following equation: \[ +1(2) + (\mathrm{Oxidation~state~of~S}) + (-2)(4) = 0 \] This yields an oxidation state of +6 for sulfur.

(f) Sulfur dioxide

Sulfur dioxide has one sulfur atom and two oxygen atoms. The chemical formula is therefore: \[ \mathrm{SO_2} \] To find the oxidation state of sulfur, we can use the rule that the sum of the oxidation states of all the atoms in a neutral molecule must be zero. The oxidation state of oxygen is -2, and there are two oxygen atoms, so the total oxidation state of the oxygens is -4. Since the sum must be zero, the oxidation state of sulfur is: \[ \mathrm{Oxidation~state~of~S} = +4 \]

(g) Mercury telluride

Mercury telluride has one mercury atom and one tellurium atom. The chemical formula is therefore: \[ \mathrm{HgTe} \] Since mercury telluride is a binary ionic compound, the oxidation state of mercury is +2. To find the oxidation state of tellurium, we can use the rule that the sum of the oxidation states of all the atoms in a neutral molecule must be zero. The oxidation state of mercury is +2. Since the sum must be zero, the oxidation state of tellurium is: \[ \mathrm{Oxidation~state~of~Te} = -2 \]

Key concepts

Chemical Formula

Understanding the chemical formula of a compound is akin to reading its unique identifier, which reveals the types and numbers of atoms present. A chemical formula is composed of element symbols from the periodic table and numerical subscriptsthat denote the amount of each element within the compound.

For instance, the formula \( \mathrm{H_2O} \) informs us that water is made up of two hydrogen atoms and one oxygen atom bonded together. In compounds like sodium thiosulfate \( \mathrm{Na_2S_2O_3} \), the formula indicates two sodium atoms, two sulfur atoms, and three oxygen atoms, which can be complex when multiple elements are involved.

Students can use this information to determine the stoichiometry of the compound—that is, the ratio in which elements combine to form the substance, which is crucial for various calculations in chemistry.

Oxidation State Rules

The oxidation state, also known as oxidation number, is a concept that provides insight into the electron distribution in chemical compounds. This is an integral part of understanding redox reactions, where electrons are transferred between species. Some general rules that govern the determination of oxidation states include:

• The oxidation state of any pure element, such as \( O_2 \) or \( N_2 \) is always zero.
• For monoatomic ions, the oxidation state is equal to the charge of the ion.
• Oxygen usually has an oxidation state of -2, except in peroxides or when bonded to fluorine.
• Hydrogen typically has an oxidation state of +1 when bonded to nonmetals and -1 when bonded to metals.
• The sum of oxidation states in a neutral molecule must be zero, while in a polyatomic ion, it must equal the ion's charge.

Applying these rules systematically allows for the determination of the oxidation state of each element in a chemical compound.

Balancing Oxidation States

Balancing oxidation states in a chemical compound is a methodical approach to ensuring that the electrons are accounted for correctly. It's a vital part of the stoichiometry in chemistry, especially during redox reactions.

For example, in sulfur dioxide \( \mathrm{SO_2} \), with sulfur having an oxidation state of +4 and each oxygen having -2, the sum of +4 and twice -2 totals zero, satisfying the requirement for a balanced molecule. It's essential to remember that the oxidation states must balance to the charge of the molecule or ion, providing a clear path for students to check their work for correctness and consistency. This also implies that the changing oxidation states in a redox reaction are interconnected and should be balanced to reflect electron transfer between species.

Polyatomic Ions

Polyatomic ions are charged species made up of two or more atoms covalently bonded together, behaving as a single unit in a chemical compound. They can be either positively or negatively charged, known respectively as cations and anions. Common examples include sulfate \( \mathrm{SO_4^{2-}} \) and ammonium \( \mathrm{NH_4^{+}} \).

Grasping the concept of polyatomic ions is pivotal in chemistry as they often appear in ionic compounds and affect the overall charge and stoichiometry of the compound. When calculating oxidation states, it's imperative to remember that the sum of oxidation states in a polyatomic ion must equal the ion's overall charge, which can sometimes lead to unusual individual oxidation states that differ from those in neutral compounds.