Question

Assign oxidation states to all of the atoms in each of the following. a. \(\mathrm{CuCl}_{2}\) b. \(\mathrm{KClO}_{3}\) c. \(\mathrm{KClO}_{4}\) d. \(\mathrm{Na}_{2} \mathrm{CO}_{3}\)

Short answer

The oxidation states for each compound are as follows: a. \(\mathrm{CuCl}_{2}\): \(\mathrm{Cu}^{2+}\mathrm{Cl}^{-}\mathrm{Cl}^{-}\) b. \(\mathrm{KClO}_{3}\): \(\mathrm{K}^{+}\mathrm{Cl}^{5+}\mathrm{O}^{-2}\mathrm{O}^{-2}\mathrm{O}^{-2}\) c. \(\mathrm{KClO}_{4}\): \(\mathrm{K}^{+}\mathrm{Cl}^{7+}\mathrm{O}^{-2}\mathrm{O}^{-2}\mathrm{O}^{-2}\mathrm{O}^{-2}\) d. \(\mathrm{Na}_{2} \mathrm{CO}_{3}\): \(\mathrm{Na}^{+}\mathrm{Na}^{+}\mathrm{C}^{4+}\mathrm{O}^{-2}\mathrm{O}^{-2}\mathrm{O}^{-2}\)

Step-by-step solution

Recall the rules for assigning oxidation states

Here are some general rules to help in assigning oxidation states: 1. The oxidation state of an atom in an elemental substance, like a single element, is 0. 2. The oxidation state of a monatomic ion equals its charge. 3. Oxygen has an oxidation state of -2 in most compounds, except for peroxides where it is -1. 4. Hydrogen has an oxidation state of +1 in most compounds, but in metal hydrides, it has an oxidation state of -1. 5. The sum of oxidation states for all atoms in a neutral compound is 0, and in an ion, it equals the charge of the ion.

Assign oxidation states for \(\mathrm{CuCl}_{2}\)

First, we will find the oxidation states for copper (Cu) and chlorine (Cl) in \(\mathrm{CuCl}_{2}\). a. In \(\mathrm{CuCl}_{2}\), the molecule is neutral, so the sum of the oxidation states of the two elements must be 0. Chlorine (Cl) is a halogen and usually has an oxidation state of -1. With two chlorine atoms, the total oxidation state from chlorine is -2. To balance this, copper (Cu) must have an oxidation state of +2. Oxidation states: \(\mathrm{Cu}^{2+}\mathrm{Cl}^{-}\mathrm{Cl}^{-}\)

Assign oxidation states for \(\mathrm{KClO}_{3}\)

Next, we will find the oxidation states for potassium (K), chlorine (Cl), and oxygen (O) in \(\mathrm{KClO}_{3}\). b. In \(\mathrm{KClO}_{3}\), the sum of the oxidation states must be 0 since the compound is neutral. Potassium (K) is an alkali metal and forms ions with a +1 charge, so its oxidation state is +1. Oxygen (O) has an oxidation state of -2. There are 3 oxygen atoms with a total oxidation state of -6. To balance this, chlorine (Cl) must have an oxidation state of +5. Oxidation states: \(\mathrm{K}^{+}\mathrm{Cl}^{5+}\mathrm{O}^{-2}\mathrm{O}^{-2}\mathrm{O}^{-2}\)

Assign oxidation states for \(\mathrm{KClO}_{4}\)

Now let's find the oxidation states for potassium (K), chlorine (Cl), and oxygen (O) in \(\mathrm{KClO}_{4}\). c. In \(\mathrm{KClO}_{4}\), the sum of the oxidation states must be 0 since the compound is neutral. Potassium (K) has an oxidation state of +1. Oxygen (O) has an oxidation state of -2. With 4 oxygen atoms, the total oxidation state for oxygen is -8. To balance this, chlorine (Cl) must have an oxidation state of +7. Oxidation states: \(\mathrm{K}^{+}\mathrm{Cl}^{7+}\mathrm{O}^{-2}\mathrm{O}^{-2}\mathrm{O}^{-2}\mathrm{O}^{-2}\)

Assign oxidation states for \(\mathrm{Na}_{2} \mathrm{CO}_{3}\)

Finally, let's find the oxidation states for sodium (Na), carbon (C), and oxygen (O) in \(\mathrm{Na}_{2} \mathrm{CO}_{3}\). d. In \(\mathrm{Na}_{2} \mathrm{CO}_{3}\), the sum of the oxidation states must be 0 since the compound is neutral. Sodium (Na) forms ions with a +1 charge, so its oxidation state is +1. With two sodium atoms, the total oxidation state for sodium is +2. Oxygen (O) has an oxidation state of -2. With 3 oxygen atoms, the total oxidation state for oxygen is -6. To balance this, carbon (C) must have an oxidation state of +4. Oxidation states: \(\mathrm{Na}^{+}\mathrm{Na}^{+}\mathrm{C}^{4+}\mathrm{O}^{-2}\mathrm{O}^{-2}\mathrm{O}^{-2}\)

Key concepts

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, involve two key processes: oxidation and reduction. These reactions are foundational in chemistry because they explain how electrons are transferred between substances.
When oxidation occurs, a substance loses electrons, leading to an increase in its oxidation state. Conversely, reduction involves the gain of electrons, resulting in a decrease in the oxidation state. In any redox reaction, one substance will donate electrons while another will accept them. Understanding these transfers is crucial as many chemical reactions, like combustion and photosynthesis, are driven by redox processes.
These reactions are characterized by changes in oxidation states, where one element's state increases and another's decreases. For instance, in water formation, hydrogen is oxidized from zero to +1, and oxygen is reduced from zero to -2, thus illustrating a redox process.

Oxidation Number Rules

Oxidation numbers are essential for determining how electrons are redistributed in reactions. Assigning them correctly requires adherence to specific rules. Understanding these rules helps in analyzing and balancing redox reactions.

• The oxidation state of any elemental substance is always zero, such as \( ext{O}_2\) or \( ext{N}_2\).
• For monatomic ions, the oxidation state equals the ion's charge. Therefore, \( ext{Na}^+\) is +1, and \( ext{Cl}^-\) is -1.
• Typically, oxygen is assigned a -2 oxidation state, although it is -1 in peroxides.
• Hydrogen usually has a +1 oxidation state, except when bonded to metals in hydrides, where it is -1.
• The sum of oxidation states in a neutral compound is always zero, while in ions, it equals the overall charge.

These rules allow for the determination of unknown oxidation states in compounds by using known values and the requirement that the sum aligns with the total charge.

Chemical Compounds Analysis

Analyzing chemical compounds for their oxidation states is a systematic process essential for interpreting redox reactions and understanding chemical behavior. This analysis incorporates applying the oxidation number rules to deduce the state of each atom within a compound.
Let's delve into a few examples to clarify. Consider \( ext{CuCl}_2\), a neutral compound. Chlorine typically presents a -1 oxidation state. With two chlorine atoms contributing a total of -2, copper must balance this with a +2 state.
Similarly, in \( ext{KClO}_3\), potassium contributes a +1 state, while oxygen, with three atoms at -2 each, provides a -6 total. To maintain neutrality, chlorine must carry a +5 state.
Through such analyses, inconsistencies in assumed oxidation states can be corrected. Careful evaluation of individual atom contributions aids in predicting the type and outcome of chemical reactions, thus forming an integral part of redox chemistry and compound analysis.