Chemical Combinations are very important in Chemistry because it aids in bringing stability to atoms of some elements especially those whose outer shells are incomplete. Atoms of elements like Neon, Helium, and Argon are completely filled with electrons as depicted in the image below while atoms of other elements are unstable but gain stability by forming chemical bonds between one another.
The cohesive force between atoms when they combine chemically is called chemical bonds. The valence electrons in an incompletely filled outer shell of an atom are called valence electrons. Only the valence electrons of an atom can take part in chemical combinations, and determine its properties. During Chemical combinations, all the valence electrons of an atom are either transferred to the valence shell of another or shared between the combining atoms in order to attain stable noble gas structures and this forms the basis of the octet rule.
The octet rule states that an atom continues to lose, gain, or share electrons in order to have eight electrons in the outermost shell.
The atoms of elements in the same group in the periodic table and isotopes of an element undergo the same type of chemical combinations, and have similar chemical properties, since they have the same number of valence electrons.
Types Of Chemical Combinations
In Chemistry, there are four main types of chemical combinations namely:
- Electrovalent(or Ionic)
- Covalent
- Co-ordinate Covalent(or Dative)
- Metallic
Electrovalent Bonding
In an electrovalent or Ionic bonding, there is a transfer of valence electrons from the donor(element with valence electrons) to the acceptor. After this transfer, both atoms gain stability and acquire a stable noble gas electronic configuration. An atom is electrically neutral but gains a charge or becomes an ion during Chemical combinations. When an atom loses electrons, it becomes a cation(positively charged) while an atom become negatively charged after accepting electrons.
An example of an Electrovalent Bonding is the bond between Sodium and Chlorine(NaCl).
Covalent Combination
In Covalent Bonding, electrons are shared between atoms of the same or different elements. Each atom contributes to the shared electrons to attain a noble gas configuration. An example of a Covalent combination is the bond between Hydrogen and Chlorine molecules.
Dative(Or Coordinate Covalent) Combination
Dative or Coordinate Covalent bonding involves the sharing of electrons however, the shared electrons are contributed by one of the participating atoms. This bond can be seen in Ammonium(NH4+).
Metallic Bonding
Metallic bonding is the process whereby the positively charged nuclei of metal atoms are simultaneously attracted to the sea of mobile electrons. A metallic bond is the electrostatic force of attraction between the positive nuclei and the sea of mobile electrons.
Intermolecular Forces
These are attractive forces that hold atoms of noble gases and covalent molecules together in liquids and solids. They are relatively weak attractive forces when compared with ionic and covalent bonds. Examples of Intermolecular forces include:
- Van der Waals Forces: Van der waals forces are weak short-range attractive forces operating between covalent molecules and they are the weakest of the intermolecular forces. They operate only when the particles are in close contact with one another as in liquids and solids.
- Dipole-Dipole(Dipolar) forces: Dipole is the partial separation of charges in a polar covalent bond and Dipole-dipole forces operate between polar covalent molecules as a result of the weak electrostatic force of attraction between the positive end of one dipole and the negative end of a neighboring dipole in polar molecules such as HCl, and HF.
- Hydrogen Bonds: A hydrogen bond is an attractive force between the hydrogen atom which is covalently bonded to an atom of flouring, oxygen, or nitrogen in a neighboring molecule.
These are the common bonds in Chemistry but more may be added in the future and this is because of the dynamic nature of science as a discipline.