Ionic or Electrovalent Bond
Table of contents
- Introduction to Ionic or Electrovalent Bond
- Formation of Ionic Bond
- Factors on Which the Formation of Ionic Bonds Depend
- Lattice Enthalpy of Ionic Crystals
- Factors on Which Lattice Enthalpy Depend
- General Properties of Ionic Bond
An atom can acquire octet by gaining the additional electrons if it has less than eight valence electrons or by losing its all valence electrons. When one or more electrons is transferred from one atom to another atom an ionic or electrovalent bond is formed. This type of bond is generally formed when a metallic atom combines with a non-metallic atom. The transfer of electrons occurs from metal atom to the non-metal atom. Thus, an electrovalent or ionic bond is formed between the atoms by the transfer of one or more valence electrons of one atom to the other atom so that they can complete their outermost octet and acquires stable nearest noble gas configuration. Therefore, in this type of bonding, one atom completes its octet and acquires noble gas configuration by losing the electrons and other by gaining the electrons.
The atom which loses the electrons acquires a positive charge and the other which gains the electrons acquires a negative charge. Because of electrostatic forces of attraction, the two oppositely charged ions come closer and form an ionic bond.
“The electrostatic force of attraction which holds the oppositely charged ions together is known as Ionic Bond or Electrovalent Bond”
The compounds containing ionic or electrovalent bonds are called Ionic or Electrovalent Compounds.
Let us illustrate the formation of ionic bond by considering the example of sodium chloride. Sodium atom (1s2 2s2 2p6 3s1) can acquire stable electronic configuration of Neon (1s22s22p6) by losing the one electron present in its valence shell.
In a similar manner like that of sodium, chlorine atom (1s22s22p63s23p5) needs only one electron to complete its octet as it has seven electrons in its valence shell. Thus, both the atoms can complete their octets if sodium atom gives one electron to chlorine atom. This tendency is responsible for bonding between sodium and chlorine atoms. Therefore, sodium gives one electron and becomes positively charged Na+ ion, while chlorine takes up the electron and becomes negatively charged Cl- ion. These two ions forms ionic bond because of electrostatic force of attraction
In terms of Lewis structure, the ionic bond between Na+ and Cl- may be represented as:
Formation of CaF2:
The electronic configurations of calcium and fluorine atoms are Ca (Z = 20) : [Ar] 4s2, F (Z = 9) : [He] 2s22p5
In the formation of calcium fluoride, calcium loses its both the valence electrons to two fluorine atoms each of which is in need of one electron. This results in the formation of Ca2+ ion and two F- ions. Each of these ions acquire noble gas configuration. One Ca2+ and two F- ions form bonds to give calcium fluoride. This may be represented as:
At the time of formation of electrovalent bond, Electrovalence or Electrovalency is the number of electrons lost or gained by an atom, it is also equal to the number of unit charges on the ion. For Example, sodium possess positive electrovalence of one, while calcium possess a positive electrovalence of two. Similarly, chlorine and fluorine both have electrovalence of one which is negative. Atoms that readily lose electrons are called electropositive while those which readily gain electrons are called Electronegative.
Similarly, formation of magnesium oxide may be shown as:
As we have learnt, ionic bond is formed by the transference of electrons between atoms. The ions, which are formed, are held together by electrostatic forces which constitute an ionic or electrostatic bond. Thus, the formation of ionic compounds depends upon:
- The ease of formation of positive and negative ions from the respective neutral atoms.
- The lattice of the crystalline compound i.e. the arrangement of positive and negative ions in a solid
The formation of positive ion involves ionization or loss of electrons while the formation of negative ion involves the addition of electrons as discussed below:
The tendency of atoms to form ionic bonds between them depends upon the following factors:
In the formation of ionic bonds, one of the atoms must form a cation by losing one or more electrons. The readiness of an atom to lose electrons depends upon its ionization enthalpy. As we know, the ionization enthalpy is the amount of energy required to remove the free electron from the outermost orbit of an isolated gaseous atom to form a positive ion.
M(g) → M+(g) + e- Ionization enthalpy
Lesser the ionization enthalpy of an atom, the greater is the ease of losing the valence electron.
Therefore, in general, metals which have low ionization enthalpy values, have greater tendency to form ionic bonds. For Example, Alkali and Alkaline Earth Metals which have relatively low ionization enthalpies, generally form ionic compounds.
The other atom participating in the formation of ionic bond must form an anion by gaining one or more electrons. This tendency of an atom to gain electrons depends upon its electron gain enthalpy. The amount of energy released when an electron is added to isolated gaseous atom to form a negative ion is electron gain enthalpy.
X(g) + e- → X- (g) Electron gain enthalpy
If the value of negative electron gain enthalpy is high, anion will be formed easily. Therefore, in general, the elements which have high gain enthalpy values form ionic compounds. For Example, halogen and oxygen group elements mostly form ionic compounds.
The formation and strength of an ionic bond also depend upon the electrostatic force of attraction between oppositely charged ions. Formation of the crystal results in release of energy due to the strong electrostatic attraction between the ions. The amount of energy released when free ions combine together to form one mole of a crystal is called Lattice Enthalpy (U),
Higher the value of lattice enthalpy greater will be the ease of formation of the ionic compound.
The ionization process is always endothermic while electron gain enthalpy process may be exothermic (releases energy) or endothermic (absorbs energy). Lattice enthalpy is exothermic (releases energy). Now, if the net effect of the above three factors is the release of energy, then Ionic Bond will be formed. This is in accordance with general observation that only those processes occur in which there is decrease of energy
Let us consider these steps for NaCl:
The ionic bond is formed because the energy released in step (ii) and (iii) is more than the energy required in step (i). As may be calculated from the above steps 640.9 kJ mol-1 of energy is released and therefore, NaCl is formed. Thus, lattice enthalpy is able to provide us with quantitative measure of the stability of an ionic compound and simply by achieving octet of electrons around the ionic species in the gaseous state does not gives full idea about stability.
Thus, to sum up, the conditions for stable ionic bonding are:
- The ionization enthalpy of atom forming the cation should be low.
- The electron gain enthalpy of atom forming the anion should be highly negative.
- Lattice enthalpy should be high.
The stability of ionic crystals is determined in terms of their lattice enthalpy. Lattice enthalpy is defined as:
“The amount of energy released when one mole of ionic crystal is formed its constituent ions in the gaseous state”.
The formation of one mole of the ionic solid from its constituent gaseous ions may be represented as:
M+(g) + X-(g) → MX(s); ∆H = Lattice enthalpy = -U
The lattice enthalpy is expressed by U. The negative sign is used because energy is released in the process. Since ionic bond is formed as a consequence of electrostatic attractions between the oppositely charged ions there is a considerable decrease of potential energy of the system. Energy needed when one mole of solid ionic compound is broken into its constituent gaseous is numerically equal to the lattice energy. Thus
MX(s) → M+(g) + X-(g) ∆H = +U
In this case, U has + ve sign because energy is absorbed in the process. Thus, lattice energy may also be defined as:
“The energy required to completely separate one mole of solid ionic compound into its gaseous ionic constituents”
It is obvious that if the compound is stable, a large amount of energy will be needed to break them. This means that higher the lattice enthalpy of an ionic compound, grater will be its stability. The magnitude of lattice enthalpy gives an idea about the interionic forces. It depends upon the following factors:
If the internuclear distance is less, then the size of the ions is also small. Consequently, the inter ionic attractions will be high and the lattice enthalpy will also be large. For Example, ionic radius of K+ (133 pm) is larger than that of Na+ (95 pm), therefore, the lattice enthalpy of NaCl (758.7 kJ mol-1) is greater than that of KCl (681.4 kJ mol-1).
Greater will be the attractive forces between the ions if magnitude of charge on the ions is large resulting in higher lattice enthalpy.
Lattice enthalpies of some ionic compounds are given in Table below.
It is evident from the Table that the lattice enthalpies of ionic solids are quite high. This is because of the strong electrostatic attractions between oppositely charged ions present in the solid. Since coulombic forces of attraction is directly proportional to the product of the charges, it follows that the higher the valency (charge) of the ions, the greater would be the lattice enthalpy of the ionic solid. Thus, lattice enthalpy increases as we move from uni-univalent ionic solids to uni-bivalent ionic solids and then to bi-bivalent ionic solids.
Most of the ionic compounds usually have cations from metallic elements and anions from non-metallic elements. But the ammonium ion NH4+ (made up of two non-metallic element) is an exception. It forms cation of many ionic compounds.
Ionic compounds usually exist, in the form of crystalline solids. X-rays studies of these compounds do not exist as independent molecules but exists in the form of ions.
These are arranged in some definite geometric patterns to form a crystal lattice For Example, in NaCl crystal each Na+ is surrounded by six Cl- ions and each Cl- is surrounded by six Na+. The number of oppositely charged ions present as the nearest neighboring around an ion is called its Coordination Number. Thus, coordination number of Na+ in NaCl ionic crystal is 6, and similarly, the coordination number of Cl- ion in NaCl crystal is 6. The crystal structure of NaCl is called Rock Salt and is shown in figure number 5
The geometric arrangement of different substances is different. It depends upon the size of the ions and the magnitude of the charges on the ions.
Ionic compounds possess high melting and boiling points, because of electrostatic forces of attractions among the ions, less amount of energy is required to break the crystal lattice. The variation in the melting point depends upon the charges on the ions and the ionic radii. The closer the ion in the crystal, the larger will be the electrostatic forces of attraction and consequently, higher will be the melting point. For Example, in case of sodium halides the melting point points decrease from NaF to NaI:
NaF (1270 K) > NaCl (1073 K) > NaBr (1023 K) > NaI (924 K)
Generally Ionic compounds are water soluble and also soluble in other polar solvents having high dielectric constants, this is because of large electrostatic interactions between ions and polar solvents. But ionic compounds are insoluble in non-polar solvents.
Ionic compounds act as good conductors of electricity when in the solution or in their molten states, as the ions become free to move. However, there is no conduction in their solid state because of strong electrostatic forces between the ions.
When the ionic compounds are dissolved in water they split up into oppositely charged ions. The chemical reactions of ionic compounds are characteristic of the constituent ions and are known as Ionic Reactions. Such reactions occur almost instantaneously.
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