Examples Of Chemical Compounds In Science
A chemical compound is, for example, a chemical substance that is formed by the bonding of two or more different chemical elements. Every chemical compound has a chemical formula, a notation that describes the relative proportions of elements that constitute a given compound. Compounds interact through chemical reactions, during which atomic bonds are broken and reform into different compounds.
The main mechanism that drives the bonding of chemical compounds is the share or transfer of valence electrons. There are three major types chemical compounds, each defined by how their atoms are bonded together: covalent compounds, ionic compounds, and metallic compounds.
Compound Vs. Mixture: Examples of Chemical Substances
A chemical compound is different than a mixture. A mixture might include a number of different elements, but in a mixture, none of the elements share chemical bonds with each other. Steel, for example, is a mixture. Steel is composed of quantities of iron (Fe) and carbon (C), but those elements are not chemically bonded. Moreover, the proportion of constituents of steel can differ. A sample of steel could have more carbon or less iron in it.
Chemical compounds, on the other hand, consist of substances that share chemical bonds with each other. Table salt (NaCl) is a compound because it is composed of a sodium atom (Na) and a chlorine atom (Cl) that are bonded together. Chemical compounds always have the same proportions of constituents; every molecule of salt is composed of one sodium and one chlorine atom (NaCl).
Types Of Chemical Compounds With Examples
There are 3 major kinds of chemical compounds, each differentiated by how the constituent atoms are held together; covalent, ionic, and metallic. What kinds of bond two elements enter into is determined by those elements electronegativities—a measure of how greatly those elements attract electrons.
Covalent compounds are chemical compounds that are formed by the sharing of electron pairs between elements. Elements will share a pair of electrons in order to gain a complete outer shell and a stable electron configuration. For example, water consists of one oxygen atom covalently bonded to two hydrogen atoms (H2O). An oxygen atom by itself only has six valence electrons; a very unstable configuration. In order to become stable, an oxygen atom will pick up two hydrogen atoms with one electron each, thus giving the oxygen atom a full shell of 8 electrons.
There are a wide variety of covalent bonds ranging from simple compounds composed of only two atoms to extremely complex organic macromolecules. Every covalently bonded compound has a specific geometric structure, determined by the intermolecular forces at play. Although most covalent bonds consist of the sharing of 2 electrons, some compounds exist where 1, 3, 4, or 6 electrons are shared.
Elements that have similar electronegativities tend to make covalent bonds. As a consequence, covalent bonds tend to be relatively weaker than the other kinds of bonds. Many common covalently bonded compounds, like carbon dioxide (CO2) or methane (CH4), exist as gasses at room temperature and their chemical bonds are relatively easy to break. That being said, some covalently bonded compounds, such as quartz or diamond, are extremely strong and have high melting and boiling points.
There exists a special type of covalent bond called a coordinate covalent bond. Coordinate covalent bonds are essentially the same in character as normal covalent bonds, except the shared pair of electrons is donated from the same element or molecule.
An ionic compound is composed of two or more ions held together by electrostatic forces. An ion forms when an atom either loses or gains an electron, thus gaining a net electrical charge. Positively charged ions are called cations and negatively charged ions are called anions. The attractive force between oppositely charged ions draws them together and forms a tight chemical bond. Individual ions in ionic compounds tend to be in close contact with a number of neighboring ions, organized in a continuous and periodic crystalline structure. Table salt, for example, is an ionic compound, formed by the joining of a sodium cation (Na+) and a chlorine anion (Cl–).
Ionic compounds generally form between elements that differ greatly in their electronegativities. For example, halogens, one of the families of elements on the periodic table, are very electronegative as they have an outer electron shell of 7 electrons. Halogens, like chlorine (Cl), readily form ionic bonds with elements that are not very electronegative, such as alkali metals like sodium (Na). Extremely electronegative elements like halogens will steal electrons from less electronegative elements, forming cations and anions. These cations and anions are then attracted to each other due to their opposite electric charges.
Ionic bonds tend to form strong, brittle compounds that are solid at room temperature, have high melting and boiling points, and are good electrical insulators. When dissolved, ionic compounds become very conductive as the constituent ions are free to move around.
Metallic compounds can be described as compounds that consist of a freely shared pool of electrons situated on a lattice of metal cations. Like ionic compounds, metallic compounds have an ordered lattice structure. Unlike ionic compounds though, the electrons of a metallic compound are delocalized, meaning that they can freely move around. The mobility of valence electrons in metallic compounds explains many of the physical properties of metals, like strength, conductivity, ductility, luster, and opacity.
Metallic bond form between metal atoms that have many open valence shells. The high number of possible energy states for electrons to occupy means that electrons can freely move about the substance. Since the electrons are very mobile, they are easily influenced by an electrical field, which makes metallic compounds very conductive. Moreover, this free-sharing of electrons means that the cation lattice can break and reform local bonds easily. This property is the reason why metals are malleable and have ductile strength. The delocalized electrons also explain why metals have their characteristic luster. Photons from light waves bounce off and are scattered by the electrons. In some ways, one could consider a sample of pure metal a single huge molecule that freely shares electrons across its structure.
It is important to realize that the above three categories only describe the three most common kinds of chemical bonds and compounds. There are other chemical compounds formed through things like hydrogen bonds, van der Waals forces, and dipole-dipole interactions. It is even theorized that massive stars could create a special type of chemical bond caused by their magnetic fields. The actual number of possible chemical bond-types is unknown, but these three categories comprise the majority of chemical bonds that one would encounter in everyday life.