Chemical bonding in analytical chemistry can be generally be separated into the three categories of ionic, covalent, and metallic bonds. Ionic and covalent bonds usually share many similarities while metallic bonds tend to produce fundamentally different molecular structures. The following is a short description and examples of each type of bond:
Ionic Bonding: these bonds occur between atoms classified as metals and non-metals, e.g. Sodium Chloride (NaCl), Potassium Bromide (KBr), and Magnesium Oxide (MgO).
Covalent Bonding: these bonds occur between two atoms classified as non-metals, e.g. Molecular Oxygen (O2), Water (H2O), and Carbon Monoxide (CO).
Metallic Bonding: these bonds occur between two atoms classified as a metal, e.g. Zinc (Zn), Iron (Fe), Metallic Brass (CuZn), and Metallic Bronze (CuSn).
PROBABLY HAVE A PERIODIC TABLE HERE THAT HAS THE METALS, SEMICONDUCTORS, AND NONMETALS COLORED DIFFERENTLY. OTHERWISE, MAYBE THIS ONE? IT IS PUBLIC DOMAIN. https://upload.wikimedia.org/wikipedia/commons/3/39/Periodic_table_large.png
The reactive nature of each atom is determined by its electric stability, this can range from the extremely reactive Fluorine (Fl) and Caesium (Cs) to Neon (Ne) and Helium (He) which have been found to only form bonds in exceedingly rare environments1. The far right column of the periodic table consists of the so called "Noble Gasses". These are defined in traditional chemistry as having exterior orbitals which have been completely filled with electrons. Most chemical reactions will consist of either the transfer of electrons from one atom to another (Ionic Bonding), the sharing of electrons between two atoms (Covalent Bonding), or the sharing fully dislocated electrons among a large mass of atoms (Metallic Bonding). The goal of these bonds is to either increase or decrease the number of electrons in the exterior orbital of an atom to the nearest full configuration. Specific examples of these can be found in their respective sections below.
Before getting deeper into each type of bonding it is important to understand the orbital theory of atoms.
Underlying each of these topics is the concept of Atomic Orbitals which can be defined as the possible energy levels the electrons are capable of occupying. These orbitals are divided into four blocks that are described below:
S Orbital: This block includes the first two columns of the periodic table from the left and Helium (He). These orbitals can contain up to two electrons a piece and generally are represented as spherical shells. Each atom in the first column of the periodic table will contain one electron it its exterior S Orbital while the second column and Helium (He) will contain two. All of these elements are classified as metals.
P Orbital: This block includes the elements in the area from Boron (B), across to Fluorine (F), and down to Oganesson (Og) on the periodic table. P Orbitals are generally described as having a “dumbbell” shape and can contain up to six electrons a piece. Each atom from column thirteen to eighteen will contain one to six electrons in its exterior orbital respectively. This region of the table is the most varied, containing metals, non-metals, and semiconductors.
D Orbital: This block contains the elements that are generally referred to as “transition” metals and spans from Scandium (Sc), across to Zinc (Zn), and down to Copernicium (Cn). Each of these D Orbitals can contain up to ten electrons but reactions in this region often include both electrons in the outer and inner orbitals. All elements within this region are metals.
F Orbital: The final region contains the so called Lanthanide and Actinide series. As can be surmised this includes both Lanthanum (La) and Actinium (Ac) and all other elements not previously covered. F Orbitals are capable of containing up to fourteen electrons. Traditionally the Lanthanide and Actinide series have been placed below the periodic table but there have been recent attempts to place it just after column two in the periodic table due to it being a more logical arrangement.
The exterior orbital of an element will generally determine the number of bonds the element will form and the structure that they will take. S Orbitals will usually only form one bond in a linear shape while P Orbitals are theoretically capable of forming up to six separate bonds. Elements in the D and F orbital regions have far less defined rules as to how bonds can form due to the very large number of electrons available in their outer shell. An atom’s orbitals will fill in a complex pattern starting with the S orbital. The exact pattern and examples of the orbital shapes can be seen in the figure below:
PROBABLY PUT EITHER THIS: https://fr.wikipedia.org/wiki/Fichier:Electron_orbitals.svg OR SOMETHING SIMILAR BELOW.
As stated above, ionic bonds will form between atoms classified as metals and atoms classified as non-metals. One or more electrons are generally transferred, also known as the electron being donated, from the metallic atoms to the nonmetallic atoms to form positively and negatively charged ions2 respectively. This variation in charge creates the electrical force that holds the bond together. An everyday example of this is the ionic bond between Sodium (Na) and Chlorine (Cl) to form table salt (NaCl). In this reaction the Sodium loses one negatively charged electron and donates it the the Chlorine atom causing each to become an ion3 Below is an explanation of the bonding process:
The Sodium donates an electron and becomes positively charged:
Na ? Na+ + e-
The Chlorine atom incorporates the electron to become negatively charged:
Cl + e- ? Cl-
The opposing charges of the Sodium and Chlorine then cause the two atoms to be drawn together forming a molecular bond. The orbital states before and after the reaction4 are given below:
[Ne]3s1 ? [Ne] + e-
[Ne]3s23p5 + e- ? [Ne]3s23p6 or [Ar]
## HERE HAVE THE DOT AND CROSS DIAGRAM FOR NaCl
Ionic molecules will by definition have an internal electrical dipole5 that can be stabilized by forming large superstructures. These superstructures are generally referred to as lattices and the exact structure is dependent upon both the individual molecular structure and the strength of the dipole. Common examples of this include table salt (NaCl)6 and Calcium Dichloride (CaCl2)7. For a sense of scale one cubic centimeter of table salt weighs roughly 2 grams. This is equivalent to roughly 2.0 x 1022 individual Sodium Chloride molecules!
The nature of ionic compounds allows dipole based fluids8, the most common of which is water (H2O) to easily break down ionic bonds and form solutions. Effectively the water and ionic molecule will form a very weak bond with the individual molecules that then allows the molecule to split. The forces between two molecules that have an electric dipole are referred to as dipole-dipole interactions9. While individual atoms have been used in these examples it is possible to form ionic molecules that replace either the positive or negative ion in the reaction.
Moving on to Covalent Bonding we begin by comparing it to the previously discussed Ionic Bonding. Rather than the donation of electrons from one atom to another, covalent bonds are usually described as sharing electrons to gain stability. As can be seen in water (H2O), one of the most common covalent molecules, covalent bonds often form in complex molecules containing more than two atoms. On the other hand, there are multiple common covalently bonded molecules that have only two atoms bonded together such as Molecular Oxygen (O2) and Molecular Nitrogen (N2).
Covalent bonds are generally divided into either a single bond or a multiple bond. The first covalent bond between two atoms will be referred to as a sigma bond (SIGMA) while the second or third covalent bonds are referred to as pi bonds (PI). Each bond will be composed of two electrons with one generally donated from each atom except in the case of coordinate bonds10 in which both electrons are donated by one of the atoms. Below is an electron diagram for some of the previously mentioned molecules:
## HAVE DOT AND CROSS DIAGRAMS HERE FOR H2, O2, H2O AS COMMON EXAMPLES
Most covalently bonded molecules will have very small electric dipoles which makes the formation of large crystalline structures uncommon11. Instead, most covalent molecules will interact with what are called intermolecular bonds. These bonds are generally much weaker than intramolecular bonds and are generally separated into the following three categories:
Permanent Dipole-Dipole Interactions: The magnitude of these forces fall in between that of the other two types. Permanent dipoles are formed when electrons are preferentially pulled closer to one atom in a bond than the other. As discussed previously, these interactions are most common in ionic molecules but can be found in some covalent molecules such as water (H2O).
van der Waals' Forces: These forces have the weakest magnitude of intermolecular forces and are often called induced dipole-dipole interactions. In analytical chemistry this is described as electrons shifting within a molecule to momentarily form electrical dipoles. These momentary dipoles will then induce dipoles in nearby molecules that can then form unstable lattices similar to those found in permanent dipole-dipole interactions. A common example of this are the pads that allow geckos to adhere to almost any surface12.
Hydrogen Bonds: These are generally the strongest of intermolecular bonds and are, as the name implies, the interactions involving a hydrogen atom and either a Oxygen (O), Fluorine (F), or Nitrogen (N) atom. Hydrogen bonds are generally seen as a specific form of dipole-dipole interactions in which the hydrogen atom has had its single electron mostly removed. This leaves the hydrogen with a largely positive electrical charge which will then be attracted to ions or the negatively charged portions of molecules. Examples of hydrogen bonding can be found in the formation of solid crystalline water (also known as ice), the mechanism by which DNA forms double helixes, and the intermolecular force that holds together cellulose fibers to create durable plant material.
As can be seen, Ionic and Covalent bonding share many similarities in how they are formed but often lead to considerably different macro structures. Later on we will return to fill in a few specifics, specifically the idea of electronegativity, that can be used to explain why some bonds are more likely than others to form.
This form of bond will almost exclusively form between metal atoms and particularly most metallic bonds will include very large numbers of metallic atoms in a lattice. Each individual metal atom will effectively remove its outer shell of electrons such that they can flow freely throughout the entire lattice. Often this delocalization of electron is referred to a "sea" of electrons. This "sea" of electrons gives rise to many of the unique properties found in metals and alloys such as their conductivity, malleability, and ductility. Descriptions of these are given below:
Conductivity: This aspect of metals is usually separated into the two categories of electrical and thermal conductivity, both of which are generally found in metallic structures. Electrical conductivity describes the ability for electrons to flow from one portion of an object to another while thermal conductivity describes the ability for vibrational energy to flow within an object.
Malleability: This quality is described as the ability of a substance to be flattened and reshaped without breaking the structure of the lattice. Most metals in their pure form are highly malleable which gives rise to the practice of creating alloys. By combining different types of metal atoms in an alloy the crystal structure can change considerably and give rise to new properties.
Ductility: This refers to the ability for a material to be stretched when force is applied. Most metallic structures have a high ductility and are easily pulled into wires.
Within a set of metallic bonds it is generally impossible to assign individual electrons to individual atoms as compared to ionic bonds. Metallic bonds also differ from covalent bonds where an electron is generally localised between two specific atoms.
As was briefly touched upon at the beginning of this article, the electrical stability of atoms tend to heavily influence the bonds they are willing to form. This variation in electrical attraction between different atoms can be measured using the Pauling Scale and is referred to as the Electronegativity of atoms. Ionic bonds generally form between atoms with large differences in Electronegativity13 whereas covalent bonds tend to form between atoms whose Electronegativity is very similar.14 The atomic Electronegativity of each element can be seen in the table below:
HERE POSSIBLY HAVE THIS PUBLIC DOMAIN ELECTRONEGATIVITY PERIODIC TABLE. https://commons.wikimedia.org/wiki/Category:Electronegativity#/media/File:Taula_peri%C3%B2dica_electronegativitat.png
Covalent bonds that have a large Electronegativity difference will produce a molecule with a non-symmetrical electrical profile that manifests in a permanent internal dipole. Molecules that have symmetrical electrical profiles will create what are called non-polar bonds while non-symmetrical ones create polar bonds. A peculiar effect of this difference can be seen in the common example of Water (H2O) and Oleic Acid15 (C18H34O2). The polar water will almost invariably separate from the non-polar oleic acid even when mixed vigorously.
As can be seen, molecular bonding imbues our universe with its readily observable properties and will generally follow a very precise set of rules. The next subjects to consider looking at in analytical chemistry would be states of matter, chemical energies, and electrochemistry.