Chemical Ionic bonding theory
Chemical Ionic bonding (metal-nonmetal):
It is formed by complete transfer of electrons from one to the other, being a metal atom and the other a nonmetal atom. Due to the atoms are neutral, when an atom gives up an electron, it is positively charged, forming what is called a positive ion or cation. The positive charge of a monovalent cation is equal in magnitude; but opposite to that of the electron (1,602 · 10-19 C) sign. If an atom captures one electron, it will be negatively charged, thereby forming a negative ion or anion. The negative charge of a monovalent anion is therefore the same as the electron. When two counterions have been formed, ie a cation and an anion, they attract each other through electrostatic type forces, and so, they can form a stable molecule. These electrostatic attractive forces, sometimes called Coulomb forces are therefore responsible for the formation of ionic compounds.
Suppose the simple case of sodium chloride or common salt. The sodium atom (Z = 11) has a single electron in orbital 3s, somewhat isolated from other pairs; while in the chlorine atom, there is also a single electron, but in this case it is in the 3pz orbital, other 3s, 3px and 3py orbitals are being inhabited by respective pairs of electrons.
The alone electron passes from sodium to chlorine, which, besides forming a pair with the electron 3pz, it will be find surrounded by other couples.
All the atoms of the alkali metals have an external configuration electronic type ns1, ie, with a single electron in the outermost orbit. This electron, which is often called valence electron is quite far from the core, which is separated also by the other electrons, called internal, which largely core-shielding attraction on said valence electron. So, It is quite easy tear this electron, for which a little energy is spent. This is the ionization energy, which, for the alkali metals, is very small.
When an alkali metal atom has been easily removed its valence electron, however, it is coster to tear a second electron, as their ionization energy is very high. Therefore, the alkali metal cations are relatively easily monovalent (M +).
In the 3 group elements atoms of the Periodic System (B, Al, Ga, In, Tl), it is cost much energy to tear valence electrons, which in this case are the type ns2np1, so it is difficult that they form trivalent cations. Heavier can form in certain cases, monovalent cations.
The elements of groups 4, 5, 6 and 7 of the Periodic Table, as well as the noble gases, the ionization energy is increasing, so it is very difficult for these elements to form positive ions. Only the heaviest elements of 4 group (tin-lead) form, in some cases, divalent cations (loss of two of the valence electrons np2). The rest of the elements of these groups form covalent bonds or negative ions.
The 7 group atoms elements of the Periodic Table (halogen) have an outer electron configuration ns2np5, that is, they lack one electron to complete the p orbitals and thus the electron configuration of noble gas that follows in the same period. Therefore, It is easy to understand that if one of these atoms are joined by a new electron, a more stable configuration is obtained, shedding energy in this process. This energy is called electron affinity, which for the halogens is high. These elements will form in a relatively easily way, monovalent anions (usually represented with the symbol X). These anions have no tendency to take a second electron, so it would have to stand, alone in an outermost orbit, without the core exercised about it any attractive force. Halogens form only monovalent anions.
Atoms of oxygen family elements (group 6 of the Periodic Table) are missing two electrons to complete the external orbitals np and acquire the noble gas configuration. Therefore, these elements tend to form divalent ions, although in this process the energy balance is slightly negative.
Nitrogen family elements (group 5 of the periodic system) are hardly trivalent anion; while carbon group (group 4) is almost impossible the form tetravalent anions. Therefore, the compounds of the nitrogen family are largely covalent part and carbon family are typically covalent.
Ionic compounds formation:
Given the sodium ionization energy and chlorine electron affinity, it can be made the energy balance, according to the equations:
Which it indicates that the formation of a pair of Cl- and Na + ions from the respective atoms in gaseous state, will spend 1.34 eV. The energy balance is even more unfavorable if one starts, as is done in the laboratory, sodium metal and molecules of chlorine gas, as in this case, must also expend the energy necessary to pass the sodium atoms from the metal into the gaseous state (heat of sublimation of sodium) and the dissociation energy to break into free atoms chlorine molecules.
In solid ionic compounds there are no molecules in the ordinary sense of the word, but rather the whole crystal can be regarded as a giant molecule, a macromolecule containing millions of ions.
In others nonionics, things happen analogously to case of sodium chloride. When anions and cations have the same loads, as, for example, in the case of calcium fluoride, comprising fluoride ions F-, and calcium ions, Ca2 +, the number of negative ions surrounding each positive ion is the twice the number of positive ions surrounding each negative ion. Thus, in the glass there is always double that number of anions of cations, which charges are compensated and is electrically neutral. Therefore, these ionic compounds represented by formulas schematically CaF2 type, although, of course, also in these cases can speak of molecules.
This cycle includes the formation of an ionic compound from the reaction of a metal (usually an element of group 1 or 2) with a nonmetal (such as halogen gases, oxygen or other) and is mainly used as means for calculating the lattice energy, it can not be determined experimentally.
The formation of gaseous ions requires the following energy processes:
- A metal atom need some energy that is called ionization energy, E to release one or more electrons.
- An atom nonmetal when it captures one or more electrons releases energy electronic -afinidad, Ae.
One might think that only ionic compounds would be formed in the event that the energy released when the nonmetal captures Electones, Ae, equals or exceeds that required for the ionization of metal, however this only happens in very few cases.
There are many ionic substances – for example sodium- chloride that are stable despite the ionization energy of the electron affinity exceeds metal nonmetal. All this makes one suspect that besides the two mentioned types of energy there are also influence by other compound formation.
The reticular energy theoretically represents the ionic compounds formation from gaseous ions. Some define it as chemical energy to break ionic compounds in gaseous ions. The first definition is exothermic and the second one is endothermic.
With the Born-Haber cycle, lattice energy is calculated by comparing the standard enthalpy of formation of the ionic compound (as elements) to the enthalpy required to make gaseous ions from the elements. This is an application of the Law of Hess.
To make gaseous ions of elements, it is necessary to convert them into gas, dissociating if it is necessary, and ionize them. If the element is a molecule (for example F2), they should be taken into account dissociation enthalpy. The energy required to remove an electron and forming a cation is the ionization energy, while the energy necessary to add and form an anion is the electron affinity.
The formation enthalpy is calculated by adding the enthalpies of atomatitation, ionization, sublimation, dissociation, electron affinity and its respective reticular energy.
Q: heat of compound formation
S: heat of metal sublimation
D: heat of non-metal dissociation
Ei: metal ionization energy
Ae: non-metal electron affinity
Er: glass energy network