Carbon is in the lowest oxidation state in (a) \(\mathrm{CH}_{4}\) (b) \(\mathrm{CCl}_{4}\) (c) \(\mathrm{CF}_{4}\) (d) \(\mathrm{CO}_{2}\)

Understanding the rules of oxidation states is crucial for analyzing chemical reactions and bonding. An oxidation state, often referred to as oxidation number, is a theoretical number assigned to an element in a chemical compound that represents its loss or gain of electrons when compared to its elemental state.

Oxidation states help in determining how electrons are distributed in a chemical compound and can indicate whether a compound is likely to be oxidized or reduced in a chemical reaction. Here are some key rules for assigning oxidation states:

• The oxidation state of a pure element is always zero.
• For monoatomic ions, the oxidation state is equal to the charge on the ion.
• Oxygen usually has an oxidation state of -2, except in peroxides where it is -1.
• Hydrogen has an oxidation state of +1 when combined with nonmetals, and -1 when bonded to metals.
• Alkali metals (group 1) always have an oxidation state of +1, and alkaline earth metals (group 2) are always +2.
• Fluorine always has an oxidation state of -1 as it is the most electronegative element.
• The sum of oxidation states in a neutral compound is zero, while in polyatomic ions it equals the charge of the ion.

These rules form a foundation for systematically assigning an oxidation state to each element within a compound.

Assigning oxidation numbers is like solving a puzzle where the oxidation rules are your guide. The goal is to find the hypothetical charges for each atom that make the overall charge of the compound come out right.

Here's a simple approach for assigning oxidation numbers: Start by assigning the elements with known oxidation states based on the rules, such as hydrogen (+1), oxygen (-2), and fluorine (-1). For a molecule like water (H2O), you'd assign hydrogen as +1 and oxygen as -2. Since there are two hydrogens, their combined charge is +2, and oxygen is -2, which equals out to zero – the overall charge for a molecule.

When dealing with molecules where the oxidation states are not immediately obvious, like in transition metal complexes, you can still apply the oxidation state rules in combination with information about the compound's charge and the known oxidation states of other ligands or atoms to deduce the oxidation state of the metal. The process requires careful analysis and sometimes a bit of trial and error to balance out the charges.

Chemical bonding is the force that holds atoms together in molecules and compounds. It encompasses a variety of interactions including ionic, covalent, and metallic bonds, all of which are influenced by the oxidation states of the participating atoms.

In ionic bonding, atoms transfer electrons to achieve noble gas electron configurations, resulting in the formation of positively and negatively charged ions. For example, when sodium (Na) combines with chlorine (Cl) to form sodium chloride (NaCl), Na gives up one electron to achieve an oxidation state of +1, while Cl accepts an electron, resulting in an oxidation state of -1.

Covalent bonding, on the other hand, involves the sharing of electron pairs between atoms, leading to the formation of molecules. The oxidation state in a covalent bond can be determined by assigning shared electrons to the more electronegative atom. For instance, in the molecule of carbon dioxide (CO2), oxygen is more electronegative than carbon, so the electrons in the C=O bonds are considered closer to oxygen, resulting in a +4 oxidation state for carbon and a -2 state for each oxygen.

Metallic bonding occurs between metal atoms, which pool their valence electrons into a 'sea' of delocalized electrons, resulting in metals' characteristic properties like conductivity and malleability. Understanding the concept of oxidation states is less directly applicable to metallic bonding, but it's still essential for grasping how metals can give up electrons to form positive ions in ionic compounds.