Complete and balance the following oxidation-reduction reactions, which give the highest possible oxidation state for the oxidized atoms. (a) \(\mathrm{K}(s)+\mathrm{H}_{2} \mathrm{O}(l) \rightarrow\) (b) \(\mathrm{Ba}(s)+\mathrm{HBr}(a q) \rightarrow\) (c) \(\operatorname{Sn}(s)+I_{2}(s) \rightarrow\)

Understanding oxidation states is essential when studying redox reactions. The oxidation state, often called oxidation number, is an indicator of the degree of oxidation of an atom in a chemical compound. Conceptually, it is the hypothetical charge that an atom would have if all bonds to atoms of different elements were fully ionic.

When assigning oxidation states, there are a few guidelines to follow. Oxygen usually has an oxidation state of -2, hydrogen is usually +1, and alkali metals in compounds have an oxidation state of +1. Transition metals can have multiple oxidation states. The sum of the oxidation states in a neutral compound must be zero, while in ions, it should equal the charge of the ion.

For the reactions given in the exercise, one would start by determining the initial oxidation states of each element. For instance, potassium (K) in its elemental form has an oxidation state of 0, while in KOH its oxidation state is +1. This change indicates a loss of one electron, which characterizes oxidation.

Balancing chemical reactions is a fundamental skill in chemistry that ensures the law of conservation of mass is adhered to. This law states that matter cannot be created or destroyed in an isolated system. To balance a chemical equation, every element must have the same number of atoms on both the reactant and product sides of the equation.

There are multiple steps involved in balancing reactions and they should be followed systematically. Begin by listing all the elements involved in the reaction and count the atoms of each on both sides. Adjust coefficients - the numbers before compounds - to achieve the same number of atoms of each element on both sides. Remember to balance elements that appear in only one reactant and one product first, and leave hydrogen and oxygen to be balanced last as they often appear in multiple compounds.

For the reactions we have, the final balanced equations take into account not only the atom count but also the charge balance, which leads into the concept of redox chemistry, involving the transfer of electrons.

Redox chemistry is the study of electron transfer in chemical reactions. Redox stands for 'reduction-oxidation'. These two processes always occur simultaneously in a redox reaction, as one species loses electrons (oxidation) while another gains electrons (reduction).

In our examples, the metals (K, Ba, and Sn) are oxidized, which means they lose electrons. Conversely, non-metals or compounds gain electrons. For example, hydrogen in H2O gains electrons to form H2 gas in reaction (a). Electron transfer can be visualized by adding electrons as reactants or products to balance the charge in the half-reactions. Understanding the electron flow allows us to explain why certain substances act as oxidizing or reducing agents.

To improve clarity for students, visuals such as oxidation number changes and electron transfer can be helpful. Additionally, discussing common oxidizing and reducing agents can provide real-world context as to how these reactions play out. Terms like 'LEO says GER' (Lose Electrons Oxidation, Gain Electrons Reduction) can serve as handy mnemonics for remembering the process.