Question

Assign oxidation states to all of the atoms in each of the following: a. \(\mathrm{Na}_{2} \mathrm{CrO}_{4}\) b. \(\mathrm{Na}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) c. \(\mathrm{Cr} \mathrm{Cl}_{3}\) d. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\)

Short answer

The oxidation states for the given compounds are: a. \(\mathrm{Na}_{2} \mathrm{CrO}_{4}\): Na = +1, Cr = +6, and O = -2. b. \(\mathrm{Na}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\): Na = +1, Cr = +6, and O = -2. c. \(\mathrm{Cr} \mathrm{Cl}_{3}\): Cr = +3, and Cl = -1. d. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\): Cr = +3, and O = -2.

Step-by-step solution

a. Assigning oxidation states for \(\mathrm{Na}_{2} \mathrm{CrO}_{4}\).

Here, the given compound is \(\mathrm{Na}_{2} \mathrm{CrO}_{4}\). 1. Sodium (Na) is an alkali metal, and its oxidation state is always +1. 2. Chromium (Cr) is a transition metal; its oxidation state can vary. 3. Oxygen (O) typically has an oxidation state of -2. Now, let's determine the charge on Cr. Since the compound is neutral, the sum of the oxidation states must be 0. 2 x (oxidation state of Na) + (oxidation state of Cr) + 4 x (oxidation state of O) = 0 2 x 1 + (oxidation state of Cr) + 4 x (-2) = 0 (oxidation state of Cr) - 6 = 0 (oxidation state of Cr) = +6 So, Na = +1, Cr = +6, and O = -2.

b. Assigning oxidation states for \(\mathrm{Na}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\).

Now, let's examine the Sodium Dichromate compound, \(\mathrm{Na}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\). 1. Sodium (Na) has an oxidation state of +1. 2. Chromium (Cr) has a variable oxidation state. 3. Oxygen (O) has an oxidation state of -2. With a neutral compound, the oxidation states' sum must also be 0. 2 x (oxidation state of Na) + 2 x (oxidation state of Cr) + 7 x (oxidation state of O) = 0 2 x 1 + 2 x (oxidation state of Cr) + 7 x (-2) = 0 2 x (oxidation state of Cr) = 12 (oxidation state of Cr) = +6 So, Na = +1, Cr = +6, and O = -2.

c. Assigning oxidation states for \(\mathrm{Cr} \mathrm{Cl}_{3}\).

Moving on to Chromium(III) Chloride compound, \(\mathrm{Cr} \mathrm{Cl}_{3}\). 1. Chromium (Cr) has a variable oxidation state. 2. Chlorine (Cl) usually has an oxidation state of -1. Here, the oxidation states' sum must also be 0 for this neutral compound. (oxidation state of Cr) + 3 x (oxidation state of Cl) = 0 (oxidation state of Cr) + 3 x (-1) = 0 (oxidation state of Cr) = +3 So, Cr = +3 and Cl = -1.

d. Assigning oxidation states for \(\mathrm{Cr}_{2} \mathrm{O}_{3}\).

Lastly, let's examine Chromium(III) Oxide compound, \(\mathrm{Cr}_{2} \mathrm{O}_{3}\). 1. Chromium (Cr) has a variable oxidation state. 2. Oxygen (O) has an oxidation state of -2. Similar to the previous examples, the compound is neutral, and the oxidation states must sum to 0. 2 x (oxidation state of Cr) + 3 x (oxidation state of O) = 0 2 x (oxidation state of Cr) + 3 x (-2) = 0 (oxidation state of Cr) = +3 Thus, Cr = +3 and O = -2.

Key concepts

Oxidation State Rules

The concept of oxidation states is a fundamental one in chemistry, providing insight into the electronic structure of atoms in compounds.
To assign oxidation states, a set of rules are followed:

• The oxidation state of a pure element is always zero.
• For ions, the oxidation state is the same as the charge.
• Elements in Group 1 of the periodic table (alkali metals) have an oxidation state of +1, and those in Group 2 (alkaline earth metals) have an oxidation state of +2.
• Oxygen usually has an oxidation state of -2, except in peroxides where it is -1, or when bonded to fluorine, where it can take a positive oxidation state.
• Hydrogen typically has an oxidation state of +1 when bonded to non-metals and -1 when bonded to metals.
• For a neutral compound, the sum of the oxidation states must equal zero.
• For a polyatomic ion, the sum of the oxidation states must equal the ion's charge.

Understanding and applying these rules helps in figuring out how electrons are distributed in chemical compounds, which is crucial for predicting reactivity and characteristics of substances.

Transition Metals Oxidation States

Transition metals are elements found in the d-block of the periodic table and are known for their ability to exist in a variety of oxidation states. This is due to the near-energy levels of the 3d and 4s orbitals.
Some key points to remember about transition metals include:

• Their oxidation states can vary widely between different compounds.
• Common oxidation states for transition metals like iron include +2 and +3, although others are also possible.
• Variable oxidation states allow transition metals to form complex ions and exhibit different properties.
• The oxidation state can influence the color of the compound in transition metals.

When determining the oxidation state of a transition metal in a compound, work with known oxidation states of other elements (like +1 for alkali metals or -2 for oxygen) and use the rule that the sum of all oxidation states must equal the charge of the molecule or ion.

Chemical Compound Neutrality

The principle of chemical compound neutrality holds that for a compound to be neutral, the total charge must be zero.
This is based on the idea that compounds are made of atoms that share or exchange electrons to reach a stable electronic configuration. For simple ionic compounds, this can be observed as the total positive charge of the cations being balanced by the total negative charge of the anions.
In polyatomic ions, while the overall ion has a net charge, the sum of the oxidation states of the atoms within that ion will equal the ion's charge.
In a neutral molecule like water, hydrogen's oxidation state is +1 and oxygen's is -2, resulting in a total oxidation state sum of zero. This concept is closely tied to the rules for assigning oxidation states and is instrumental in writing correct chemical formulas and equations as well as in understanding the structure and reactivity of molecules.