For our purposes at this point in the text, we can define an acid as a substance with at least one hydrogen atom that can dissociate to form an anion and an H⁺ ion (a proton) in aqueous solution, thereby forming an acidic solution. We can define bases as compounds that produce hydroxide ions (OH⁻) and a cation when dissolved in water, thus forming a basic solution. Solutions that are neither basic nor acidic are neutral. Pure acids and bases and their concentrated aqueous solutions are commonly encountered in the laboratory. They are usually highly corrosive, so they must be handled with care.

Acids

The names of acids differentiate between (1) acids in which the H⁺ ion is attached to an oxygen atom of a polyatomic anion (these are called oxoacids, or occasionally oxyacids) and (2) acids in which the H⁺ ion is attached to some other element. In the latter case, the name of the acid begins with hydro– and ends in –ic, with the root of the name of the other element or ion in between. Recall that the name of the anion derived from this kind of acid always ends in –ide. Thus hydrogen chloride (HCl) gas dissolves in water to form hydrochloric acid (which contains H⁺ and Cl⁻ ions), hydrogen cyanide (HCN) gas forms hydrocyanic acid (which contains H⁺ and CN⁻ ions). Examples of this kind of acid are commonly encountered and very important. For instance, your stomach contains a dilute solution of hydrochloric acid to help digest food. When the mechanisms that prevent the stomach from digesting itself malfunction, the acid destroys the lining of the stomach and an ulcer forms.

If an acid contains one or more H⁺ ions attached to oxygen, it is a derivative of one of the common oxoanions, such as sulfate (SO₄²⁻) or nitrate (NO₃⁻). These acids contain as many H⁺ ions as are necessary to balance the negative charge on the anion, resulting in a neutral species such as H₂SO₄ and HNO₃.

The Relationship between the Names of the Oxoacids and the Names of the Parent Oxoanions

Bases

We will present more comprehensive definitions of bases in later chapters, but virtually every base you encounter in the meantime will be an ionic compound, such as sodium hydroxide (NaOH) and barium hydroxide [Ba(OH)₂], that contain the hydroxide ion and a metal cation. These have the general formula M(OH)_n. It is important to recognize that alcohols, with the general formula ROH, are covalent compounds, not ionic compounds; consequently, they do not dissociate in water to form a basic solution (containing OH⁻ ions). When a base reacts with any of the acids we have discussed, it accepts a proton (H⁺). For example, the hydroxide ion (OH⁻) accepts a proton to form H₂O. Thus bases are also referred to as proton acceptors.

Concentrated aqueous solutions of ammonia (NH₃) contain significant amounts of the hydroxide ion, even though the dissolved substance is not primarily ammonium hydroxide (NH₄OH) as is often stated on the label. Thus aqueous ammonia solution is also a common base. Replacing a hydrogen atom of NH₃ with an alkyl group results in an amine (RNH₂), which is also a base. Amines have pungent odors—for example, methylamine (CH₃NH₂) is one of the compounds responsible for the foul odor associated with spoiled fish. The physiological importance of amines is suggested in the word vitamin, which is derived from the phrase vital amines. The word was coined to describe dietary substances that were effective at preventing scurvy, rickets, and other diseases because these substances were assumed to be amines. Subsequently, some vitamins have indeed been confirmed to be amines.

In this concept, the arrangement of bonds around the central atom is considered to depend upon the number of valence shell electron pairs, and on the relative sizes and shapes of these orbitals. These arrangements hold good for non-transition elements i.e., those elements which do not use electrons in bond formation. The geometrical shapes are actually result of the tendency of the electron-pairs to remain at a maximum distance apart so that the interaction between them is minimum. The repulsion between free electron-pairs will be obviously greater than that of repulsion between a bond pair and another bond pair. Let us sum up the essential features of this theory under the following rules:

1. The preferred arrangement of a given number of electron pairs in the valence shell is that which makes them to remain at a maximum distance apart.
2. A non-bonding pair (lone pair) of electrons is capable of making more space on the surface of an atom than a bonding pair. This is because the non-bonding electron pair is under the influence of one nucleus only but the bonding electron pair is constrained by two nuclei.
3. The influence of a bonding electron pair decreases with the increasing value of electronegativity of an atom forming a molecule.
4. The two electron pairs of a double ,bond (or the three electron pairs of a triple bond) take up more space than the one electron pair of a single bond.

Applications of Valence Shell Electron Pair Repulsion Concept:

Let us now apply the valence shell electron pair repulsion concept to explain the shapes of the molecules. In other words, the effect of electron pair repulsion on molecular structure will be discussed. The shapes of the molecules and ions of non-transition elements will now be described in terms of this theory. The molecules will be classified according to the number of electron pairs present in them, irrespective of the fact whether they are of bonding or non-bonding type.

The molecialar orbital theory describes the valence electrons as associated with all the nuclei concerned. The nuclei are in equilibrium positions in the stable molecule and electrons associated with all the nuclei can be described by wave functions. The energy states of electrons can be described in the combined states or molecular orbitals. The molecular orbitals are multicentred or delocalised. They are filled with the required number of electrons (each molecular orbital is usually filled with two electrons). Molecular orbitals may be obtained by the linear combination of atomic orbitals (LCAO method).

The molecular orbitals are assumed to possess the following characteristics:

(i) Each electron in the molecule is described by a wave function The value of (1) is such that the value of 1)2 at any point represents the probability of finding the electrons in unit volume around that point. The wave functions are called molecular orbitals. These molecular orbitals are polycentric so that the electron moves in the field of all the nuclei.

(ii) Each molecular orbital has its own energy. 

(iii) Each electron has a definite spin \+ 2 or — —2i and Pauli’s exclusion principle is observed.

(iv) The appropriate form of the wave equation is quite complicated and cannot be used for exact solution except for hydrogen. Thus approximations are necessary. One of the approximations is that when an electron comes in the vicinity of one nucleus, the force arising on it is due to the nucleus and its other electrons. Both the wave equation and its solutions resemble those for the isolated atom, and the molecular orbital consists of a series of superposed self-consistent orbitals. This procedure is known as the linear combination of atomic orbitals (LCA0).

(v) The greater the overlap of atomic orbitals among themselves, more stable molecular orbitals (with least energy states) are obtained.

(vi) The energy of a molecular orbital is least when the combining atomic orbitals have equal or almost equal energy states. Atomic orbitals of low energy will not be able to overlap with other atomic orbitals and electrons carried by them will be non-bonding.

(vii) Each molecular wave function corresponds to a definite energy value. The sum of the individual energies of the molecular orbitals, after correction, represents the total energy of the molecule.

Let us now apply these factors to a simple homonuclear diatomic molecule such as hydrogen in which two identical atoms are linked by an electron pair bond.

The covalent bonds are formed from the overlap of the orbitals in the region between the centres of two atoms. As a result of it the nuclei of bonded atoms approach each other more closely than do the nuclei of non-bonded atoms. The covalent radius of an atom is taken as one half the distance between the nuclei of two identical atoms forming a single covalent bond.

Thus we can say that the bond distance between the two atoms A — B should be the arithmetic mean of the bond lengths A — A and B — B. Take the case of covalent radius of carbon which is one half the experimentally determined distance between C — C single bond. This gives the value of 77 pm. Similarly, for the Si — Si linkage the covalent radius comes to be 117 pm. Now if we consider the bond distance between carbon and silicon, we should expect a bond length of 194 pm. This is in very good agreement with the experimentally determined C — Si distance of 193 pm in carborundum (silicon carbide). Covalent radius of C — Si bond is half of the bond length.

The covalent radii decrease with increase in bond order because there is corresponding decrease in internuclear distance. Thus in carbon, the internuclear distance and atomic radii decrease with increase in bond.

Bond Order Inter Nuclear Distance (pm) Atomic Radius (pm) C – C 154 77 C=C  134 67

C =- C 120 60

Although the above rule works in many simple diatomic molecule, but this is not generally the case. Very often there is a considerable deviation from the expected result. This deviation can be attributed to many factors like multiple.

Mendeleev’s Periodic Table in spite of its advantages suffers from the following drawbacks and has thus limitations in its application:

Position of hydrogen is not clear because it resembles with both alkali metals and halogens. It gives the positive H+ ions like alkali metals and gives the hydride ions like halides.

Certain chemically similar elements, e.g., copper, gold, platinum are placed in different groups while some dissimilar elements are grouped together.

Certain elements of higher atomic weight precede others with lower atomic weight. Argon (At. Wt. = 40) Precedes Potassium (At. Wt. = 39) Cobalt (At. Wt. = 59) Precedes Nickel (At. Wt. = 58.6)

No position is assigned to isotopes in different groups.

Maximum valence state is depicted by an element in a particular group. The elements of group VIII usually do not depict 8 maximum oxidation state except ruthenium and osmium.

No explanation is available for the inert pair effect and stability of valence states differing by units of two.

Mendeleev interpreted this relation between elements on the basis of Periodic Law which he stated as:

“The properties of elements are periodic functions of their atomic weights.”

However, certain discrepancies were noted while arranging the elements on the basis of atomic weights. For example, beryllium was out of place in the table as its atomic mass was 13.5 which should fit in between carbon and nitrogen. Similarly, inert gases had no proper place in the Periodic Table.

These problems were not fully resolved until 1914 when Moseley showed that elements could be arranged in a periodic pattern on the basis of their atomic numbers. The Periodic Law now states as:

“The properties of elements are periodic functions of their atomic numbers.”

By taking atomic numbers as the basis of the periodic classification of’ elements, various anomalies and misfits have been removed. For example, proper positions to cobalt and nickel, potassium and argon, etc., have been given.

The classification of elements was an interesting arrangement and attracted considerable attention. Several gaps in the table suggested discovery of new elements. Chemical and physical properties of unknown elements could be predicted which helped in the search of new elements. Mendeleev suggested that elements similar to aluminium and silicon should exist. Gallium, Ga (similar to aluminium) was discovered in 1875 and Germanium, Ge in 1886.

Mendeleev solved the problem of odd elements (which could not be adjusted properly in a group) by dividing groups into subgroups ‘A’ and ‘W. For example, among metals of the first group, sodium and potassium were placed in Group IA and copper and silver in Group 1B because of the difference in properties. Noble gases were discovered in the last decade of the nineteenth century and were placed in zero group because of their inertness.

The main features of Mendeleev’s Periodic Table were the arruigement of elements in vertical columns or groups and the horizontal rows or periods. He left spaces for the unknown elements and predicted properties of the Germanium which was not discovered until 1886. He called it eka-Silicon as it fell below Silicon.

In order to understand the Soft and Hard Acid-Base concept, it is essential to know the meanings of Lewis Acids and Lewis Bases. A Lewis base is a lone pair electron donor and a Lewis acid is a lone pair electron acceptor.

When a Lewis acid (E) combines with a Lewis base (N), a chemical bond results e.g.,

                                                 E+.N —> E:N or E  N

When a pair of electrons is held by a cr bond between two different atoms which differ widely in size, electro negativity etc., the bonding pair will be held more tightly to one core than to the other. A bond of this type is generally highly polar and relatively labile and is referred to as coordinate bond.

When the rates of reactions are considered, the Lewis acids are called Electrophiles and Lewis bases are known as Nucleophiles. The Lewis acids include most of the cations while the Lewis bases are mostly the anions and neutral spacies. If we break an organic molecule conceptually, we see that it is also a combination of Lewis acid and a Lewis base e.g., C2H5OH, where C2115+ is a Lewis acid and OH- is a Lewis base. Hence all carbonium ions (although may not exist freely) are considered to act as Lewis acids (Electrophiles), since they contain such a structure which can accept a pair of electrons from the Lewis base. Similarly, OH- ions act as Lewis base (Nucleophile).

Classification of Acceptor and Donor Atoms and Ions:

In 1958, Chatt and Coworkers divided Lewis Acids (acceptor molecules and ions) into two classes:

Class (a): Those Lewis acids which form their most stable complexes with the first member of Group V, VI & VII in the Periodic Table i.e., N, 0, F (which act as donor atoms or ligands).

Class (b): Those Lewis acids which form their most stable complexes with the donor atoms (ligands) of the subsequent elements of these groups i.e., P, S, 13r, etc.

The donor atoms and ions (Lewis Bases) were classified on the basis of electronic affinity, coordinating ability, effective charge, ionic size and polarization considerations.

The electron affinity sequences of various groups of electron pair donor atoms and ions (ligands) with respect to the class (a) and class (b) electron pair acceptors (Lewis acids) is given below:

Class (a)

F >> Cl > Br > I

O>> S > Se > Te

N << P > As > Sb  

Class (b)

F Cl < Br < I

0 << S — Se — Te

N >> > As > Sb

It is observed that greater the values of electron affinities between donor-acceptor atoms or ions greater will be their coordinating affinities. Thus, in general more stable complexes of donor atoms i.e., F, 0,. N, etc., will be formed with class (a) acceptors and class (b) acceptors (Lewis Acids) will form less stable complexes with F, 0, N in their respective oxidation states. Polarization of the donors (ligands) by the acceptOr also plays an important role in determining the stabilities of the complexes.

Based on the polarization considerations, Pearson introduced the idea of HARD and SOFT acids and bases. According to him, the Lewis bases (ligands) which are more polarizable are ‘Soft’, and those which are less polarizable are `Hard’. For example, the. atoms F, 0, and N are the hardest Lewis bases. Hence, Pearson’s concept of Hard and Soft acids and bases is in close agreement with class (a) and (b) acceptors given by Chat and Coworkers, Class (a) refers to hard acids and class (b) to soft acids.

Pearson, based on the concept of polarizability, divided the Lewis acids and bases as defined below:.

Hard Bases:

The donor atoms of low polarizabilities, high electronegativities and associated with empty orbitals of high energy are classed as hard bases. They are hard to oxidise.

Soft Bases:

The donor atoms of high polarizabilities, low electronegativities and associated with empty orbitals of low energy are termed as soft bases. They are easy to oxidize.

Hard Acids:

These are acceptor atoms of high positive charge, small ,size and do not have outer electrons which can be easily excited.

Soft Acids:

These are acceptor atoms of low positive charge, large size and have several outer electrons which can be easily excited.

Based on these considerations, Pearson classified the Lewis acids and Lewis bases as hard and soft as given below:

Classification of Lewis Bases:

Hard                                                                                  Soft

H20, OH-, F-                                                              R2 S, RSH, RS –

CH3C0i, PO43, SO42                                               1-, SCN-, S2032

C1-, CO32, C10:1, NO3                                              ; R3P, R3As, (R0)3P

NH3, RNH2, N2H4                                                    C2H4 C6H6 H- R-                         

                                              Border line

                              C6H5NH2, C5H5N, N3, Br-, NO.;, S032, N2

A brief review of the various concepts regarding acids and bases are given here.

THE ARRHENIUS (CLASSICAL) CONCEPT (1884)

According to this an acid is defined as a hydrogen containing substance which gives 114 ions (i.e., H30+ hydronium ions) when dissolved in water.

A base is a substance which coliains OH groups and gives hydroxyl ions OH – when dissolved in water.

Arrhenius concept is based upon ionic dissociation of compound in water. For example, HCl is an acid because it produces H30+ ions in water but CH4 is not. Similarly, NaOH is a base because it furnishes OH- ions, whereas C2H5OH is not a base.

                                       HC1 + H2O  > H30+ + Cl-

                                       NaOH + H2O  > Na+ + OH

The process of neutralization of an acid by a base can be represented by the reaction to form neutral water.

H+ + OH  > H2O

Advantages:

With this concept, many aspects of acid-base behavior were understood. For instance, the constant heat of neutralization of a strong acid by a strong base can readily be explained in terms of Arrhenius concept because the reaction ‘involves only the combination of a hydrogen ion and a hydroxyl ion in all such neutralization reactions.

It explains the catalytic properties of acids. Arrhenius theory affords a correlation between the electrolytic dissociation and the concentrations of the hydrogen ion. The mobility of the hydrogen ions parallel the catalytic activity of the solution if the hydrogen ion is truly the source of the catalytic properties.

Shortcomings:

According to this theory, all the acid-base reactions are limited to aqueous medium only. It does not explain the acid-base reactions taking place in non-aqueous solvents such as liquid ammonia.

It also cannot explain the reactions in gas phase where no solvent is present.

Similarly, the definition of a base under this concept is restricted to compounds containing hydroxyl ions only, whereas many organic compounds as well as ammonia which exhibit basic properties cannot be explained by this definition. Similarly, there are many acidic compounds which do not contain hydrogen and cannot be explained on the basis of Arrhenius concept. Hence new concepts were put forward to explain more general cases of acids and bases.

THE PROTONIC OR LOWRY-BRONSTED CONCEPT (1923)

According to Bronsted:

An acid is defined as a species (a compound or an ion) which donates or tends to donate a proton (II+ ion).

A base is a species which accepts or tends to accept a proton. Acid-base reaction is the transfer of a proton from an acid to a base. The dissociation of an acid HA can be represented as:

HA  > A- + H+

Acid Base proton CH3COOH  > CH3C00- +h

Acid Base Proton

According to this definition, any negatively charged ion (anion) acts as a base. Thus, CH3C00- is a base and is said to be conjugate base of acetic acid In an acid-base reaction, an acid yields a base (conjugate) and base after accepting proton yields a conjugate acid.

The acid-base reaction is represented as:

Al Ba  > B1 A2

Bronsted Conjugate acid base & base acid

The conjugate acid-base pairs are species on opposite sides of an equation that differ by a proton. The weaker acids have stronger conjugate base pairs and stronger acids have weaker conjugate base.

Thus Cl S024 , OH – are conjugate bases of HC1, HSO4 and H2O, respectively. Similarly, H2O, HSO4 and HC1 are conjugate acids of the bases OW, S024 and Cl -, respectively.

The following species may be regarded as acids:

Molecular Species: HC1, H2SO4, CH3COOH, HCN, H2S, H2O etc.

Anionic Species:HSO4, HCO3, H2PO4, HP024, HS etc.

Cationic Species: H30+, NHS, [Cu(H20)4]2+ etc.

The following species may be regarded as bases:

Molecular Species: H2O, NH3, CH3NH2 etc.

Anionic Species: OH-, HS -, S2-, HCO3, HSO4, C1- etc.

Cationic Species: [Fe(H20)50H]2+, [Cu(H20)3 Of1]+ etc.

From the above examples, it is found that some of the species act both as acids and bases depending upon the manner they behave in the given reaction. Amphiprotic Species: A species that acts bath as a proton donor and a proton acceptor is said to be amphiprotic. For example:

(i) H2O is amphiprotic. It loses proton to a base such as NH3 or accepts a proton from an acid such as HC1.

H20 + NH3  > NH+4 + OH-

Acid Base

H20 + HCl  > H30+ + Cl-

Base Acid

The proton-containing negative ions are amphiprotic. For example:

HS- + OH –  > S-2 + H2O

Acid Base

HS- + H30+  > H2S + H2O

Base Acid

The conjugate base and acid are shown as:

OH- Conjugate base of water

— H+ H2O +H+

H+ HCO3

H30+ Conjugate acid of water CO3 Conjugate base of HCO3

H2CO3 Conjugate acid of HCO3

Amphoteric hydroxides react with both acids and bases becadse they are equally amphiprotic.

Al(H20)3 (OH)3 + OH –  > Al(H20)2 (OH)4 + H2O

Acid Base

Al(H20)3 (OH)3 + H2O+  > Al(H20)4 (OH)2 + H2O

Base Acid

Polyprotic Acids: Acids containing one proton which can be donated are monoprotic acids. Those acids which contain more than one donatable proton are known as “Polyprotic acids” e.g., HC1, HNO3, HCN etc. are monoprotic, whereas H3PO4, H2SO4, H3AsO4 are Polyprotic.

Non-aqueous solutions also lose or gain protons and fit into the Bronsted acids and bases. For example, water is dissolved in liquid ammonia.

H20 + NH3  > OH- + NH+4

Acid Base

Certain elements exist in more than one oxidation state and thus show variable valency. For example, transition metals show wide range of valence states, e.g., Fe2+, Fe3+; Cr2+, Cr3+ Co2+, CO3+, (more common), Co4+ and Co5+ (less common) etc. Similarly, normal metals, e.g., Pb2+ and Pb4+ ; Sn2+ and Sn4+ also show variable valencies. Certain non-metals are also found to show more than one valence states e.g., P3+ and P5+.

Variable valency may be due to different reasons and would be discussed over here accordingly.

In transition elements (elements in which d orbitals are in the process of completion) the variable valency is due to the partially filled ‘ d’ orbitals, state of hybridization and type of reactants (usually called ligands). Such elements can involve different number of electrons in compound formation and would show variable valencies. The stability of the particular valence state would also depend upon the nature of the reacting species and the number of d electrons present.

Iron (Fe) shows normally Fe2+ and Fe3+ valence states. Atomic number of Fe is 26 and its electronic configuration would be: 1s2 2s2 2p6 3s2 3p6 3d6 4s2.

Loss of two electrons from 4s orbitals would not leave behind an inert gas configuration. Therefore, more chances of electrons to be pulled out exist. Since `cl’ orbitals would be more stable when half-filled (with d5 configuration), so loss of one electron from 3d orbital along with two electrons from 4s orbitals would give a more stable state of the ion. Therefore, iron in Fe3+ state would be more stable. Let us show the valence shell having 3d orbitals and arrange the electrons in accordance with valence bond theory (See Chapter 4). In order to get next inert gas configuration more electrons would be required which are supplied by reacting species (Lewis bases or ligands) and the ions are thus stabilized. Mostly water will be acting as Lewis base and due to this reason most of the transition metal salts exist in stable state as hydrated species.

Elements of lanthanides and actinides show variable valencies due to the involvement of ‘f’ orbitals.

Some of the `1,’ block elements show variable valency due to the involvement of an ‘inert pair’ of electrons. The ‘inert pair’ of electrons are present in the ‘s’ orbitals and do not take part in chemical reactions under the prevailing conditions. Under such conditions only `p’ orbitals would take part in bond formation. However, if the ‘inert pair’ of electrons present in `s’ orbitals is actuated to take part in bond formation, an increase in valence state by two units will be observed. It is due to the `inert pair’ of electrons that the difference in valence states of such elements is always by two units.

Although the inert pair effect is quite marked for bivalent tin compounds but tetravalent tin is more stable of the two. Hence, bivalent tin is readily converted to tetravalent state, and thus the former ion is a reducing agent.

Sn(II) — 2e- Sn(IV)

The `13′ block elements which have vacant ‘cl’ orbitals available in their electronic configuration also show variable valency by utilizing these orbitals in addition to the corresponding ‘s’ and `ps’ orbitals. Thus the involvement of vacant ‘cr orbitals would be responsible for the variation in valence state of such elements. Let us consider the first two members of Vth group namely, nitrogen and phosphorus. There is no chance for the presence and involvement of `cr ‘orbitals in nitrogen as indicated by its electronic configuration (1s2 2s2 2p3) i.e., no ‘ar orbital is available in the 2nd shell.

This type of bond is exhibited by atoms which can either lose electrons to form positively charged ions (cations) or gain electrons to form negatively charged ions (anions). The atom which can lose electrons is said to be electropositive or basic and the atom capable of gaining one or more electrons is referred. as electronegative. The more electropositive atom has always low value of ionization potential and is thus capable of losing electrons with greater ease. The electrons lost by electropositive atoms are completely transferred to other atoms which show greater electro negativities. The bond formed by complete transfer of electrons from electropositive atom to more electronegative atom is called ionic or electrovalent bond The electropositive elements in energy terms should have higher energy states than those of electronegative elements. This ‘energy difference will be responsible for the flow of electrons from higher energy states to lower energy states. The two atoms are held together by electrostatic forces of attraction acting between such atoms.

M° + X°                 M+ : X (3.1)

The energy required to completely separate the ions from a diatomic molecule is given by the following expression:

Potential energy = Electrostatic energy + Van der Waals’ energy.

The electrostatic energy — q1 q2 e  R

where q1 and q2 are charges on atoms M and X and R is the internuclear separation.

The general tendency of various atoms to form molecules is to attain inert gas configuration, being the most stable. The atoms of the inert gases have outermost p orbitals completely filled. Such a configuration will not easily lose or gain any electron because very high ionization potentials (or electro negativities) will be required to remove an electron or gain any additional electron. Let us consider potassium and chlorine atoms which would combine to form potassium chloride molecule. Potassium atom (Atomic number of K = 19) has the electronic configuration 1s2 2s2 2p6 3s2 3p6 4 • si. Chlorine (Atomic number of Cl = 17) has electronic configuration 1s2 2s2 2p6 3s2 3p5. None of these has an inert gas structure. But they possess an incomplete shell of electrons and orbitals. The nearest inert gas to both is argon having electronic configuration 1s2 2s2 2/ 3s23p6. Thu:s the lois of one electron from potassium and gain of this lost electron by chlorine would leave both the atoms with argon configuration. During this process, IC+ and would be produced and ‘the electrostatic attraction between these oppositely charged Ions should be responsible for a stable ionic bond.

Let us take some more examples to elaborate the ionic bonding situation.

• Formation of Sodium Chloride Molecule
• Formation of Magnesium Chloride Molecule
• Formation of Aluminium Fluoride

Transition metal ions which do not resort to inert gas configuration attain their stability through the formation of complex ions. The ions with odd number of s or p electrons are not known, but an odd number of d electrons is found in transition metal ions.

It should be noted that important forces between atoms or groups of atoms are electrostatic in nature.