Introduction

Atoms are the basic building blocks of matter. When multiple atoms interact, they can form molecules. The type of bond that forms, either covalent, ionic, or metallic, depends on how the bond’s electrons interact.

In chemistry, everything tends to move toward stability, and atoms are no exception. They bond with one another to reach a more stable, lower-energy state.

Valence Electrons

Valence electrons, not core electrons, are typically the ones involved in bonding. They play a key role in how atoms bond and largely determine the chemical properties of the molecules they form.

Electronegativity

Before discussing bonds, you should be familiar with the key periodic trend of electronegativity. Here are the essentials to keep in mind:

Coulomb's Law

Coulomb's Law is useful for understanding the attraction between two charged particles, like atoms in a bond. It tells us that the strength of the electrostatic force holding them together depends on and is proportional to two main factors.
Here is the equation: $F=k\frac{q_1{q}_2}{r^2}$

This ties directly into the trend of electronegativity and actually helps explain it. As you move down a group on the periodic table, the distance between the valence electrons and the nucleus increases. This weaker attraction results in a lower electronegativity. When electrons are farther from the nucleus, the pull just isn’t as strong.

When comparing the melting points of ionic compounds, look for ions with higher magnitudes of charges and smaller sizes, since stronger electrostatic attractions require more energy to break, resulting in higher melting and boiling points.

Ionic Bonding

Ionic bonds form through the transfer of valence electrons from one atom to another, typically from a metal to a nonmetal. One example is the formation of sodium chloride.

$2Na(\mathrm{s})+{\mathrm{Cl}}_2(\mathrm{g})\to 2NaCl(\mathrm{s})$

In this reaction, sodium donates one electron to chlorine, forming ${\text{Na}}^{+}$ and ${\text{Cl}}^{-}$ ions. These oppositely charged ions are held together not by shared electrons but by strong electrostatic forces, which are the attraction between positive and negative charges. This strong attraction is why ionic compounds like $\text{NaCl}$ have high melting and boiling points.

Ionic compounds also form a crystal lattice structure, where ions are arranged in a repeating three-dimensional pattern. This rigid structure contributes to their brittleness and strength. 

When ionic compounds are melted or dissolved in water, the ions are free to move, allowing the substance to conduct electricity. That’s what makes ionic compounds such good conductors in liquid or aqueous form.

Covalent Bonding

In covalent bonds, which usually form between nonmetals, electrons are shared between atoms, and whether the bond is polar or nonpolar depends on the difference in electronegativity; polar covalent bonds have an unequal distribution of charge, while nonpolar covalent bonds share electrons equally.

To understand the difference between polar and nonpolar covalent bonds, we need to look at electronegativity. When atoms with similar electronegativity share valence electrons, they form a nonpolar covalent bond, like two oxygen atoms sharing electrons equally because they pull with the same strength. Think nonpolar = balanced.

On the other hand, when atoms with unequal abilities to attract electrons form a bond, a polar covalent bond forms. For example, in a water molecule, oxygen is farther to the right on the periodic table compared to hydrogen, which means it has a stronger pull on electrons. This causes the shared electrons to be drawn closer to oxygen, creating an uneven charge distribution with oxygen gaining a partial negative charge and hydrogen a partial positive charge.

This difference leads to the formation of bond dipoles, and the greater the electronegativity difference, the stronger the dipole. The bigger the electronegativity difference, the stronger the bond dipole.

$\begin{array}{cc}\text{Bond}& \text{Electronegativity Difference}\\ \text{Nonpolar covalent}& <0.4\\ \text{Polar covalent}& 0.4-1.8\\ \text{Ionic}& >1.8\end{array}$

This chart gives a rough idea of the electronegativity differences that help determine the nature of a bond. You don’t need to memorize the exact values, but having a general sense of them can be helpful when figuring out whether a bond is ionic, polar covalent, or nonpolar covalent.

Ionic Bonds form between two elements with a significant electronegativity difference between a metal and a nonmetal. They result from the attraction between a positively charged ion (cation) and a negatively charged ion (anion).

Covalent bonds usually form between two nonmetals.

Characteristics of Ionic and Covalent Compounds.

Metallic Bonding

Metallic solids are made up of metal atoms arranged in a closely packed structure. What makes them unique is the presence of delocalized electrons, electrons that are not tied to any one atom and can move freely throughout the entire structure. This is often described as a “sea of electrons” surrounding positively charged metal ions.