Write the complete Lewis structure for (a) ammonium chloride; (b) potassium phosphide; (c) sodium hypochlorite

At the heart of chemical bonding, lies the concept of valence electrons. These are the electrons that reside in the outermost shell of an atom and are pivotal in determining an atom's capacity to form bonds with others. In simpler words, think of valence electrons as the hands of an atom, reaching out to hold onto other atoms in a chemical handshake. For example, in ammonium chloride (NH₄Cl), nitrogen's 'hands' are full with its own electrons plus holding hands with four hydrogen atoms.

Understanding the number of valence electrons gives insight into how many bonds an atom can form. This concept is not only foundational for creating Lewis structures but also for understanding the stability of molecules. Elements in the same group of the periodic table have the same number of valence electrons, which leads to similarities in their bonding characteristics.

Ionic bonds are like a game of give and take between atoms. One atom, typically a metal, will give up its valence electrons to another, often a non-metal, which is looking to gain electrons to feel complete. This creates positively charged ions (cations) and negatively charged ions (anions) that are attracted to each other by their opposite charges, just as the opposite ends of magnets are drawn together.

For instance, in potassium phosphide (K₃P), potassium (the giver) hands over its one valence electron to phosphorus (the taker), which needs three more to be satisfied. This electron exchange results in a strong ionic bond between K⁺ and P^3-. Understanding ionic bonding helps you to predict the properties of the substance, such as its strength, melting point, and whether it will conduct electricity when dissolved in water.

The way electrons are distributed among atoms in a molecule is called electron distribution. This concept is central to drawing Lewis structures, as it directly depicts which atoms are sharing electrons in covalent bonds or transferring electrons in ionic bonds. Electron distribution helps in visualizing molecule geometry and understanding how atoms are linked.

Let's consider sodium hypochlorite (NaOCl). Sodium (Na) parts with its lone valence electron, contributing to the electron needs of chlorine (Cl) and oxygen (O), which share electrons between them. This sharing and transferring create a complete electron 'map' of the molecule, showing all individual and shared electrons, thus giving insight into the actual bonding nature of the compound. Keeping a tally of valence electrons and following the octet rule for main-group elements are key practices to accurately distribute electrons in Lewis structures.

Chemical bonding is the strong force that holds atoms together in molecules. There are mainly two types of bonds: covalent and ionic. Covalent bonds involve the sharing of electron pairs between atoms, while ionic bonds result from the attraction between positively and negatively charged ions. Through these bonds, atoms attain stability, usually by achieving a full outer electron shell. That could either mean sharing electrons to fill up the valence shell, as in a water molecule (H₂O), or transferring electrons to empty or fill the valence shell, as seen in the formation of sodium chloride (NaCl).

Understanding chemical bonding is crucial for predicting how a substance will interact with others, its reactivity, phase of matter, and even its conductivity. The type of bond greatly affects the properties of the molecules formed and, thus, the behavior of the substance under different conditions.