Chemical bonding

• Chemical bond is the attractive force holding atoms or ions together.
• This attractive interaction leads to a more stable state for the whole system compared to individual atoms.
• Valence electrons play a fundamental role in chemical bonding.
• In the electron configuration of an atom, the outermost shell is called the valence shell, and the electrons in the valence shell (outermost shell) are known as valence electrons.
• The electron configuration of carbon is 1s 2 2s 2 2p 2.
• The outermost shell of carbon is the 2nd principal shell, so there are 4 valence electrons in carbon.
• Valence electrons are the electrons that are the furthest away from the nucleus, and thus they experience the least attraction from the nucleus and are the most reactive.
• There are two major types of chemical bonds: ionic bonds and covalent bonds.
• An ionic bond is a bond that results from the electrostatic attraction (force) between ions of opposite charges.
• Ionic bonds apply to ionic compounds, such as sodium chloride (NaCl).
• In simple ionic compounds, the metal element loses valence electron(s) to form the cation and the non- metal element gains electron(s) to form the anion.
• According to Lewis’s Theory, an atom is most stable if its outer shell is filled or contains eight electrons.
• When one atom loses an electron and another atom gains that electron, the process is called electron transfer.
• Sodium (Na) only has one electron in its outer electron shell, so it is easier (more energetically favorable) for sodium to donate that one electron than to find seven more electrons to fill the outer shell.
• Chlorine (Cl), on the other hand, has seven electrons in its outer shell.
• In this case, it is easier for chlorine to gain one electron than to lose seven, so it tends to take on an electron and become Cl -.
• A covalent bond is a bond formed through the sharing of electron pairs between the two bonding atoms.
• The shared electron pairs are mutually attracted by the nuclei of both atoms.
• By sharing the electron pairs, both atoms also gain a filled outer shell, or an octet.
• Almost all the bonds involved in organic compounds are covalent bonds.
• Covalent bonds can be non-polar or polar.
• For covalent bonds formed between two identical atoms, the electron pairs are shared equally between the two nuclei.
• Electron density is distributed evenlythrough the bond, making it a non-polar bond.
• Examples of covalent bonding include all homonuclear molecules, such as H-H, Cl-Cl, O=O, N≡N.
• A single water molecule, H 2 O consists of two hydrogen atoms bonded to one oxygen atom.
• LDF arise from the formation of temporary instantaneous polarities across a molecule from the circulations of electrons.
• The term “INTERmolecular forces” is used to describe the forces of attraction between atoms, molecules, and ions when they are placed close to each other.
• The dipole-dipole interactions that result from these dipoles is known as hydrogen bonding.
• When two particles experience an intermolecular force, a positive (+) charge on one particle is attracted to the negative (-) on the other particles.
• As a result, molecules with higher molecular weights have higher LDF and consequently have higher melting points, boiling points and enthalpies of vaporization.
• Hydrogen bonding is an especially strong form of dipole-dipole interaction.
• The relative values of EN are listed using the scale devised by Linus Pauling, as summarized in the following table:
• Intermolecular forces are more likely to be found in condensed states such as liquid or solid.
• When a hydrogen atom is covalently bonded to nitrogen, oxygen or fluorine a very strong dipole is formed.
• All pure metallic elements consist of metal bonds, for example, gold (Au), iron (Fe), aluminum (Al), etc.
• When intermolecular forces are weak, the atoms, molecules or ions do not have a strong attraction for each other and move far apart.
• An instantaneous polarity in one molecule may induce an opposing polarity in an adjacent molecule, resulting in a series of attractive forces among neighboring molecules.
• The polarities of individual molecules tend to align by opposites, drawing molecules together and thereby favoring a condensed phase.
• The greater the ΔEN, the more polar the bond is.
• Metallic bonds are a strong kind of bond that is spread out like a network.