Question

Assign oxidation states to all of the atoms in each of the following: a. \(\mathrm{CuCl}_{2}\) b. \(\mathrm{Cr} \mathrm{Cl}_{3}\) c. \(\mathrm{HCrO}_{4}^{-}\) d. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\)

Short answer

The oxidation states for the given compounds are: a. \(\mathrm{CuCl}_{2}\): Cu: +2, Cl: -1 b. \(\mathrm{CrCl}_{3}\): Cr: +3, Cl: -1 c. \(\mathrm{HCrO}_{4}^{-}\): H: +1, Cr: +6, O: -2 d. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\): Cr: +3, O: -2

Step-by-step solution

Assign oxidation state to Cl

Since Cl is a monoatomic ion (it is in the form of a single chlorine atom), its oxidation state is equal to its charge, which is -1.

Determine the oxidation state of Cu

The sum of the oxidation states in the compound must equal the overall charge of the compound, which is 0 in this case; thus, we have: Cu + 2(-1) = 0 Cu - 2 = 0 Cu = +2 So, the oxidation state of Cu is +2. Hence, the oxidation states in \(\mathrm{CuCl}_{2}\) are: Cu: +2 Cl: -1 #b. \(\mathrm{CrCl}_{3}\)#

Assign oxidation state to Cl

As mentioned before, the oxidation state of Cl is -1.

Determine the oxidation state of Cr

The sum of the oxidation states in the compound must equal the overall charge of the compound, which is 0 in this case; thus, we have: Cr + 3(-1) = 0 Cr - 3 = 0 Cr = +3 So, the oxidation state of Cr is +3. Hence, the oxidation states in \(\mathrm{CrCl}_{3}\) are: Cr: +3 Cl: -1 #c. \(\mathrm{HCrO}_{4}^{-}\)#

Assign oxidation states to H and O

The oxidation state of H is +1, and the oxidation state of O is -2.

Determine the oxidation state of Cr

The sum of the oxidation states in the ion must equal the overall charge of the ion, which is -1 in this case; thus, we have: H + Cr + 4(-2) = -1 1 + Cr - 8 = -1 Cr = +6 So, the oxidation state of Cr is +6. Hence, the oxidation states in \(\mathrm{HCrO}_{4}^{-}\) are: H: +1 Cr: +6 O: -2 #d. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\)#

Assign oxidation state to O

The oxidation state of O is -2.

Determine the oxidation state of Cr

The sum of the oxidation states in the compound must equal the overall charge of the compound, which is 0 in this case; thus, we have: 2(Cr) + 3(-2) = 0 2Cr - 6 = 0 Cr = +3 So, the oxidation state of Cr is +3. Hence, the oxidation states in \(\mathrm{Cr}_{2} \mathrm{O}_{3}\) are: Cr: +3 O: -2

Key concepts

Redox Reactions

Redox reactions are chemical processes where the oxidation states of atoms are changed. These reactions involve the transfer of electrons between substances. One substance undergoes oxidation, losing electrons, while the other experiences reduction, gaining electrons. When a substance is oxidized, its oxidation number increases. Conversely, during reduction, the oxidation number decreases. It’s like balancing a seesaw, where one side rises as the other falls. In practical terms, the oxidizing agent gains electrons and the reducing agent loses electrons. Understanding redox reactions is crucial because they are fundamental to many natural and industrial processes such as cellular respiration and corrosion. By identifying which elements have undergone changes in their oxidation states, chemists can decipher the specific roles elements play in reactions.

Oxidation Numbers

Oxidation numbers, also known as oxidation states, serve as helpful tools in tracking how many electrons an atom has gained, lost, or shared. Think of them as an accounting system for electrons in compounds. Each element in the periodic table comes with a set of possible oxidation states it can attain. For instance:

• Hydrogen usually has an oxidation state of +1.
• Oxygen generally carries an oxidation state of -2.
• Halogens like chlorine often have an oxidation state of -1.

Assigning these numbers involves considering the overall charge of the compound and applying specific rules. This helps in determining which element has been oxidized and which has been reduced in a reaction. Through these numbers, chemists ensure that the process of redox reactions remains balanced.

Chemical Compounds

Chemical compounds are substances composed of two or more different atoms bonded together. Each compound has distinct chemical properties arising from the types of atoms and the manner in which they are bonded. In compounds, elements behave differently than they do in their elemental forms. For instance, sodium chloride (table salt) is a compound formed by sodium and chlorine, which individually have quite different properties. Compounds are characterized by their molecular formulas, which indicate the elements involved and their ratios. Understanding these formulas is crucial as it allows us to determine the oxidation states of each element. This, in turn, provides insight into the compound's stability and reactivity. With the aid of rules and conventions, we can interpret these formulas to understand the behavior and characteristics of specific chemical compounds.

Oxidation State Rules

Determining oxidation states is guided by a set of fundamental rules. These rules ensure consistency when analyzing the electron distribution in compounds. Here's a concise rundown to help you figure things out:

• Elements in their pure state always have an oxidation state of zero.
• For a monoatomic ion, the oxidation state equals the ion's charge. For example, Cl^- has an oxidation state of -1.
• In molecules, the sum of all oxidation states must match the overall charge of the compound or ion.
• Oxygen typically has an oxidation state of -2, while hydrogen typically has +1.

These rules guide chemists in consistently determining which atoms in a compound are oxidized and which are reduced. Keep these principles in mind, and you’ll find it easier to navigate through complex chemical equations and recognize redox reactions. They are the foundation upon which much of chemistry understanding is built.