Question

Assign oxidation states to all of the atoms in each of the following. a. \(\mathrm{CBr}_{4}\) b. \(\mathrm{HClO}_{4}\) c. \(\mathrm{K}_{3} \mathrm{PO}_{4}\) d. \(\mathrm{N}_{2} \mathrm{O}\)

Short answer

The oxidation states for the given compounds are as follows: a. CBr\(_4\): C: +4 Br: -1 b. HClO\(_4\): H: +1 Cl: +7 O: -2 c. K\(_3\)PO\(_4\): K: +1 P: +5 O: -2 d. N\(_2\)O: N: +1 O: -2

Step-by-step solution

a. Assign oxidation states for atoms in \(\mathrm{CBr}_{4}\)

In this molecule, we only need to focus on carbon (C) and bromine (Br). The oxidation state of carbon is not known, so let's call it x. There are four bromine atoms, each with an oxidation state of -1 (since bromine is a halogen). According to rule 3, the sum of all oxidation states in a neutral compound is zero: x + 4(-1) = 0 Solving for x, we get: x = 4 So the oxidation states are: C: +4 Br: -1

b. Assign oxidation states for atoms in \(\mathrm{HClO}_{4}\)

In this compound, we have hydrogen (H), chlorine (Cl), and oxygen (O). The oxidation states are: - H: +1 (because it is bonded to a more electronegative atom, oxygen) - Cl: let's call it x (unknown) - O: -2 (since it is bonded to a less electronegative atom, chlorine) Now we can use rule 3 and set up an equation for this neutral molecule: +1 + x + 4(-2) = 0 x - 7 = 0 Solving for x, we get: x = 7 So the oxidation states are: H: +1 Cl: +7 O: -2

c. Assign oxidation states for atoms in \(\mathrm{K}_{3} \mathrm{PO}_{4}\)

This compound consists of a cation, \(\mathrm{K}^{+}\), and an anion, \(\mathrm{PO}_{4}^{3-}\). For potassium (K), it's a monatomic ion with a charge of +1. For the polyatomic ion \(\mathrm{PO}_{4}^{3-}\), we have phosphorus (P) and oxygen (O). The oxidation states are: - P: let's call it x (unknown) - O: -2 Now we can use rule 4 to set up an equation for the polyatomic ion: \( x + 4(-2) = -3 \) Solving for x, we get: x = +5 So the oxidation states are: K: +1 P: +5 O: -2

d. Assign oxidation states for atoms in \(\mathrm{N}_{2} \mathrm{O}\)

In this compound, we have nitrogen (N) and oxygen (O). The oxidation states are: - N: let's call it x (unknown); since there are two nitrogen atoms, their total oxidation state is 2x. - O: -2 Now we can use rule 3 to set up an equation for this neutral molecule: 2x + (-2) = 0 Solving for x, we get: x = +1 So the oxidation states are: N: +1 O: -2

Key concepts

Redox Reactions

Redox reactions are chemical reactions where oxidation and reduction occur simultaneously. These processes always happen together, which is why they are called "redox" (short for "reduction-oxidation"). In every redox reaction, one element loses electrons, undergoing oxidation, while another gains electrons, undergoing reduction. Let's break it down with an example.

Consider the interaction between sodium (\( ext{Na}\)) and chlorine (\( ext{Cl}_2\)) to form sodium chloride (\( ext{NaCl}\)). In this reaction, sodium atoms lose electrons and become \( ext{Na}^+\) ions (oxidation), while chlorine molecules gain electrons to become chloride ions (\( ext{Cl}^-\)) (reduction).

To identify redox reactions:

• Look for changes in oxidation states of elements involved in the reaction.
• Identify the element being oxidized, which will have an increase in the oxidation state.
• Identify the element being reduced, which will have a decrease in oxidation state.

Understanding these changes is crucial as redox reactions play vital roles in many biological and industrial processes, such as metabolism and combustion.

Oxidation Number Rules

In chemistry, oxidation numbers (or oxidation states) are used to track electron transfer in chemical reactions. Assigning oxidation numbers can help to determine what happens in a reaction. There are established rules to follow when determining oxidation numbers:

• The oxidation state of an atom in a pure element is always 0. For example, Ne in its natural state is 0.
• For monatomic ions, the oxidation state is equal to the charge of the ion. For instance, in \( ext{Na}^+\), the oxidation state is +1.
• In compounds, hydrogen generally has an oxidation state of +1, and oxygen has an oxidation state of -2 (with some exceptions).
• The sum of oxidation states in a neutral compound must be zero, but in a polyatomic ion, it equals the ion's charge.

These rules help balance reactions and verify the transfer of electrons, ensuring that during redox reactions, the total increase in oxidation states equals the total decrease.

Chemical Bonding

Chemical bonding involves the joining together of two or more atoms to form molecules. Bonds form because atoms seek to reach a lower, more stable energy state. There are several types of chemical bonds, with covalent and ionic being the most common.

In **covalent bonding**, electrons are shared between atoms. This type of bond typically occurs between two non-metals. For example, in water (\( ext{H}_2 ext{O}\)), each hydrogen atom shares electrons with the oxygen atom.

**Ionic bonding**, on the other hand, involves the transfer of electrons from one atom to another, leading to the formation of ions. This occurs between metals and non-metals. For instance, in sodium chloride (\( ext{NaCl}\)), the sodium atom donates an electron to the chlorine atom, resulting in positively charged sodium ions and negatively charged chloride ions, which attract each other.

The nature of these bonds determines the physical and chemical properties of substances, such as melting points, solubility, and conductivity.

Oxidizing and Reducing Agents

In redox reactions, substances that facilitate oxidation or reduction are known as oxidizing or reducing agents. These agents play complementary roles in redox reactions:

An **oxidizing agent** is a substance that accepts electrons and is reduced in a chemical reaction. Basically, it "oxidizes" another substance by gaining electrons. For example, in the reaction of hydrogen peroxide (\( ext{H}_2 ext{O}_2\)) as an antiseptic, it acts as an oxidizing agent by accepting electrons.

A **reducing agent** donates electrons, becoming oxidized itself, and reduces another substance. For example, hydrogen gas (\( ext{H}_2\)) can act as a reducing agent by donating electrons, as seen in its reaction with fluorine to form hydrogen fluoride (\( ext{HF}\)).

Understanding the nature and roles of these agents:

• Helps in identifying how substances interact in various chemical processes.
• Supports recognizing which conditions favor oxidation or reduction.
• Aids in controlling chemical reactions in industrial applications, like metallurgy or battery technology.

Essentially, these agents are crucial for driving the transformations seen in redox reactions.