Arsonic and nitric acids react to form nitrogen monoxide, arsenic acid and water according to the equation shown below. Is this an example of a redox reaction? \(2 \mathrm{HNO}_{3}(a q)+3 \mathrm{H}_{3} \mathrm{AsO}_{3}(a q) \rightarrow 2 \mathrm{NO}(g)+3 \mathrm{H}_{3} \mathrm{AsO}_{4}(a q)+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})\)

Understanding oxidation states is crucial when analyzing redox reactions, which involve the transfer of electrons between chemical species. An oxidation state, often referred to as an oxidation number, is a theoretical charge that an atom would have if all bonds to atoms of different elements were completely ionic.

The rules for assigning oxidation states include considering known values, such as the fact that oxygen usually carries a -2 charge (except in peroxides or when bonded to fluorine), hydrogen has a +1 when bonded to non-metals, and metals typically have positive values corresponding to their group number in the Periodic Table.

Accurately determining oxidation states is the key first step in identifying redox processes. Once these states are known, changes in them between reactants and products reveal which atoms are oxidized and which are reduced.

Chemical reactions are processes that transform one set of chemical substances into another. They are the foundation of chemistry and involve the making and breaking of chemical bonds.

Reactions can be categorized into many types, with redox reactions being one such type. A redox reaction consists of two concurrent processes: oxidation, where a substance loses electrons, and reduction, where a substance gains electrons. These reactions are ubiquitous in both nature and technology, contributing to processes such as metabolism and battery operation.

Balancing chemical equations is necessary to satisfy the law of conservation of mass, which states that matter cannot be created nor destroyed in a chemical reaction. Thus, the quantity of each element must be the same in both the reactants and products.

To balance a chemical equation, one should adjust the coefficients in front of compounds until the number of atoms for each element is the same on both sides of the reaction. For redox reactions, it's often helpful to balance the equations by ensuring that the number of electrons lost in oxidation equals the number of electrons gained in reduction.

Reduction and oxidation are the two key processes that comprise redox reactions. Oxidation is the loss of electrons, which increases an atom's oxidation state, while reduction is the gain of electrons, decreasing an atom's oxidation state.

These processes are interconnected: whenever one substance is oxidized, another is reduced. This reciprocity is known as the oxidation-reduction pairing and is represented in an equation where the oxidizing agent gets reduced, and the reducing agent gets oxidized.

An easy way to remember this is the mnemonic 'OIL RIG'—which stands for 'Oxidation Is Loss, Reduction Is Gain' of electrons. In the exercise example, arsenic is oxidized (it becomes more positive), and nitrogen is reduced (it becomes more negative), demonstrating a classic redox reaction.