Question

Assign oxidation states to all of the atoms in each of the following: a. ${\mathrm{NF}}_3$ b. ${\mathrm{NO}}_2$ c. $\mathrm{HOCl}$ d. $\mathrm{S}\${\$}_8$

Short answer

In summary, the oxidation states for each compound are as follows: a. ${\mathrm{NF}}_3$: N: +3, F: -1 b. ${\mathrm{NO}}_2$: N: +4, O: -2 c. $\mathrm{HOCl}$: H: +1, O: -2, Cl: +1 d. ${\mathrm{S}}_8$: S: 0

Step-by-step solution

Part a: ${\mathrm{NF}}_3$

First, write down the usual oxidation states of the elements involved: 1. Nitrogen (N) typically has oxidation states of -3, +1, +3, or +5. 2. Fluorine (F) almost always has an oxidation state of -1. Since we have 3 Fluorine atoms with an oxidation state of -1 each, we can calculate the oxidation state of Nitrogen as follows: $N=-3(F1+F2+F3)=-3(-1+(-1)+(-1))=-3(3)=+3$ Therefore, the oxidation states in ${\mathrm{NF}}_3$ are: - Nitrogen (N): +3 - Fluorine (F): -1

Part b: ${\mathrm{NO}}_2$

In this case, the usual oxidation states of the elements are: 1. Nitrogen (N) typically has oxidation states of -3, +1, +3, or +5. 2. Oxygen (O) typically has an oxidation state of -2. As there are 2 oxygen atoms with an oxidation state of -2 each, we can calculate the oxidation state of Nitrogen: $N=-2(O1+O2)=-2(-2+(-2))=-2(-4)=+4$ Thus, the oxidation states in ${\mathrm{NO}}_2$ are: - Nitrogen (N): +4 - Oxygen (O): -2

Part c: $\mathrm{HOCl}$

The usual oxidation states of the elements are: 1. Hydrogen (H) typically has an oxidation state of +1. 2. Oxygen (O) typically has an oxidation state of -2. 3. Chlorine (Cl) typically has oxidation states of -1, +1, +3, +5, or +7. We can calculate the oxidation state of Chlorine using the given oxidation states for Hydrogen and Oxygen: $H+O+Cl=0$ $+1+(-2)+Cl=0$ $Cl=1$ The oxidation states in $\mathrm{HOCl}$ are: - Hydrogen (H): +1 - Oxygen (O): -2 - Chlorine (Cl): +1

Part d: ${\mathrm{S}}_8$

In a molecule consisting of only one type of element, like ${\mathrm{S}}_8$, the oxidation state of each atom in the molecule is zero. This is because the atoms have no charge or are not bonded to a more (or less) electronegative atom. The oxidation state of each Sulfur atom in ${\mathrm{S}}_8$ is: - Sulfur (S): 0

Key concepts

Chemical Compounds

Chemical compounds are substances that consist of two or more different types of atoms bonded together. These atoms are combined in fixed ratios and have specific properties that differ from the individual elements that make them up. For instance, Nitrogen trifluoride (NF₃) is a compound made of nitrogen and fluorine atoms. While individual nitrogen and fluorine atoms have their own unique characteristics, when combined, they create a stable compound with distinct properties.
In the world of chemistry, it's essential to understand how atoms combine to form compounds. These combinations happen because of chemical bonds, which can be ionic or covalent. In ionic compounds, atoms transfer electrons to one another, whereas in covalent compounds, such as those mentioned in the exercise— 2) and Hypochlorous acid (HOCl)> —the atoms share electrons.

• Ionic bonds occur when electrons are transferred, forming ions.
• Covalent bonds occur when electrons are shared between atoms.

Understanding chemical compounds is important because it helps us know how elements interact with each other. This interaction determines the chemical and physical properties of a compound, such as its reactivity, color, state, and toxicity.

Oxidation State Calculations

The process of calculating oxidation states allows us to determine the electron distribution in a compound. This helps us understand the chemical behavior of substances. Each element in a compound is assigned an oxidation state, which reflects its ability to lose or gain electrons . Oxidation state calculations are grounded in a few basic rules:

• The oxidation state of a free element (an element not combined with another in a compound) is zero. This is why sulfur in S₈ is zero.
• For binary compounds, the element that is more electronegative is assigned a negative oxidation state corresponding to its typical ion form.
• Fluorine always has an oxidation state of -1 in compounds 3>.
• Oxygen typically has an oxidation state of -2, except in peroxides.
• Hydrogen usually has an oxidation state of +1, except when it forms hydrides with metals.

By applying these rules, we can deduce the oxidation state of each atom in compounds. For instance, in NO₂, where nitrogen is combined with oxygen, the calculations show how nitrogen maintains an oxidation state of +4 due to the presence of two oxygen atoms each with -2.

Elemental Oxidation States

Elemental oxidation states are fundamental concepts in chemistry that denote the hypothetical charge an atom would have if all its bonds to different atoms were completely ionic. This concept is particularly important when analyzing redox reactions, as it helps identify which elements are oxidized and which are reduced. For example, in the compound HOCl, hydrogen has an oxidation state of +1, oxygen is -2, and chlorine is +1. Here, each element's oxidation state helps us understand their roles in chemical reactions.
Sometimes, elements can have multiple oxidation states, as seen with chlorine which can vary from -1 up to +7. This versatility is typically influenced by the particular compound the element is in and the specific atoms it combines with. Understanding elemental oxidation states involves recognizing how different elements naturally tend to gain, lose, or share electrons:

• Highly electronegative elements, like fluorine, almost always have negative oxidation states.
• Metals often have positive oxidation states as they tend to lose electrons.
• Noble gases generally remain at an oxidation state of zero due to their lack of tendency to engage in chemical bonding.

Thus, mastering this concept is key to predicting and balancing chemical equations and understanding the nature of substances involved in reactions.