Question

Which of the following species has an atom with \(+6\) oxidation state? (a) \(\mathrm{MnO}_{4}^{-}\) (b) \(\mathrm{Cr}(\mathrm{CN})_{6}^{3-}\) (c) \(\mathrm{NiF}_{6}^{2-}\) (d) \(\mathrm{CrO}_{2} \mathrm{Cl}_{2}\)

Short answer

The species with an atom in a \(+6\) oxidation state is (d) \(\mathrm{CrO}_{2} \mathrm{Cl}_{2}\).

Step-by-step solution

Understand the concept of oxidation states

The oxidation state of an element in a compound represents the number of electrons that an atom has gained or lost compared to its elemental state. In a neutral compound, the sum of the oxidation states of all elements must equal zero. In a polyatomic ion, the sum of the oxidation states must equal the charge of the ion.

Analyze option (a) - \(\mathrm{MnO}_{4}^{-}\)

Oxygen typically has an oxidation state of \-2. Since there are 4 oxygen atoms, their total oxidation state will be \(4(-2) = -8\). To balance the charge of the permanganate ion (\(\mathrm{MnO}_{4}^{-}\)), the oxidation state of Mn must be \(\text{+7}\) to have \(\text{{+7}} + (-8) = -1\), which matches the charge of the ion.

Analyze option (b) - \(\mathrm{Cr}(\mathrm{CN})_{6}^{3-}\)

Cyanide (CN) carries a \-1 oxidation state. Since there are 6 cyanides, their total oxidation state is \(6(-1) = -6\). To balance the \(3-\) charge of the ion, chromium must have a \(+3\) oxidation state, so \(\text{{+3}} + (-6) = -3\), equating to the charge of the ion.

Analyze option (c) - \(\mathrm{NiF}_{6}^{2-}\)

Fluorine has an oxidation state of \-1. With 6 fluorine atoms, their combined oxidation state is \(6(-1) = -6\). The nickel ion must balance the \(2-\) charge with a \(+4\) oxidation state since \(\text{{+4}} + (-6) = -2\), equalling the charge of the ion.

Analyze option (d) - \(\mathrm{CrO}_{2} \mathrm{Cl}_{2}\)

Each chlorine in \(\mathrm{CrO}_{2} \mathrm{Cl}_{2}\) has an oxidation state of \-1. Therefore, the two chlorines contribute \(2(-1) = -2\) in total. Each oxygen has an oxidation state of \-2, so for two oxygen atoms, it's \(2(-2) = -4\). To balance the compound, which is not an ion and thus has an overall neutral charge, the oxidation state of Cr must be \(\text{{+6}}\) as \(\text{{+6}} + (-4 + -2) = 0\).

Key concepts

Redox Chemistry

Redox chemistry, short for reduction-oxidation chemistry, is a fundamental concept that deals with the transfer of electrons between substances. It's an essential part of understanding how chemical reactions occur, especially those that involve the gain or loss of electrons.

In simple terms, oxidation refers to the loss of electrons, leading to an increase in oxidation state, while reduction refers to the gain of electrons, leading to a decrease in oxidation state. In any redox reaction, one species gets oxidized as the other gets reduced; this is called a complementary process. For instance, in a battery, redox reactions are what allow for the flow of electrons from one end to the other, powering the devices it's connected to.

The significance of redox chemistry extends beyond just theoretical aspects; it's instrumental in industries such as energy (in the form of batteries and fuel cells), metallurgy (in the extraction of metals), and biology (in cellular respiration and photosynthesis).

Determination of Oxidation Number

Determining the oxidation state, or oxidation number, of an element within a chemical species is a key skill in redox chemistry. The oxidation state is an indicator of the degree of oxidation (loss of electrons) or reduction (gain of electrons) the atom has undergone in comparison to its elemental form.

Several rules help us assign oxidation numbers:

• The oxidation state of a pure element is always zero.
• For ions composed of only one atom, the oxidation state is equal to the net charge.
• Oxygen is usually assigned an oxidation state of -2, with exceptions in compounds like peroxides.
• Hydrogen is typically +1 when bonded with nonmetals and -1 with metals.
• The sum of oxidation states in a neutral molecule equals zero. In a polyatomic ion, the sum must equal the ion’s charge.

The correct identification of oxidation states is crucial for analyzing redox reactions, as observed in the given exercise, where comparing oxidation states determines the answer.

Chemical Bonding

Chemical bonding is the force that holds atoms together in compounds, creating stability in the arrangement of electrons. There are several types of chemical bonds, including ionic, covalent, and metallic bonds.

Ionic Bonds

Ionic bonds form when one atom gives up one or more electrons to another atom, resulting in the formation of ions. These ions are held together by electrostatic forces. For example, in sodium chloride (NaCl), sodium (Na) loses an electron and becomes positively charged, while chlorine (Cl) gains that electron to become negatively charged.

Covalent Bonds

Covalent bonds involve the sharing of electrons between atoms. This sharing allows each atom to achieve a full outer shell of electrons, according to the octet rule. Water (H₂O), with two hydrogen atoms each sharing their single electron with oxygen to fill its outer shell, is a prime example.

Metallic Bonds

In metallic bonds, electrons are not shared between individual atoms. Instead, they are delocalized throughout the entire metal structure, which explains the conductivity and malleability of metals.

Understanding chemical bonding is essential for analyzing how atoms combine and how this affects the physical and chemical properties of substances.