structure-reactivity relationship

Introduction

In this chapter, we will learn about the consequence of molecular structure on the reactivity of the molecules. The reactivity of 1 material towards another is calculated through the rate at which the two substances react and the amount of the products shaped.

Not all molecules are equally reactive. But, what build several organic molecules more reactive than others. To discover an answer to this question, we should have various thought of the nature of reactions that the organic molecules undergo. A huge number of reactions that the organic molecules undergo can be readily understood as easy analogies of acid-base reactions. Consequently, it is significant for us to know the basic traits of acid-base reactions. We will start this chapter through a discussion on diverse methods in which the acids and bases can be described. We will then familiarize us through the idea of acid-base equilibrium. Here, we will as well learn that the position of the acid-base equilibrium is a compute of molecular reactivity; additional it is influenced through many factors. Even though, the functional groups present in a molecule are of key significance in determining the molecular reactivity, it has been examined that several compounds containing similar functional groups differ in their reactivities. Therefore, additionally to the occurrence of the functional groups the nature and arrangement of atoms connected to the functional groups as well manage the molecular reactivity. Such consequences which are connected through the transform in molecular structure are termed in this chapter; we will learn several structural consequences these as inductive result; resonance effect and steric effect, and their influence on molecular reactivity.

In addition to the structural consequences, we will as well talk about solvent effects and hydrogen bonding that is as well significant factors affecting the rate and the extent of such reactions. At last, we will learn an interesting symmetry involving a proton shift from 1 atom of a molecule to another, termed tautomerism.

Acids and Bases:

There are several methods of defining acids and bases. According to Arrhenius (1884), a Swedish chemist, an acid is a material which ionizes in aqueous solution to create hydrogen ions (H⁺), as well identified as protons. And, a base is a material that ionizes to create hydroxide (OH) ions. Therefore, Arrheniun theory supposes an easy dissociation these as,

That during dissociation, the covalent bond between H - A is broken and the electrons shaping this bond shift on A as given via the curved arrow.

Therefore, HCl is an acid and NaOH is a base because on dissociation they yield H⁺ and OH^- ions, correspondingly. Hence, the strength of such acids and bases is related to the degree of their dissociation. The mineral acids such as HCl, Hl, HBr, H₂SO₄ and HNO₃ are tough acids since they are roughly entirely dissociated in aqueous solutions. Likewise, the strength of a base will as well base upon its degree of dissociation.

An alternative theory of acids and bases was devised independently via Bronsted and Lowry in the year 1922. According to the Bronstzed-Lowry approach, an acid is a proton donor and a base is a proton acceptor. Because under normal reaction states a free proton can't subsist as a divide entity, whenever an acid in the Bronsted-Lowry sense is considered, a base must be present to admit the proton from the acid. The Bronsted acids are as well termed protic acids since they react via the relocate of a proton. Consider the following example, here; the ethanoic acid is an acid because it donates a proton to water which is a base because it accepts the proton. Similarly, the ethanoate ion, which is formed by the loss of a proton from ethanoic acid, functions as a base since it can admit a proton to become ethanoci acid again. Therefore, ethanoate ion is termed the conjugate base of ethanoic acid. Likewise, the hydronium ion is the conjugate acid of the base, water. This pair of a base and its conjugate acid or an acid and its conjugate base is as well termed conjugate acid-base pair.

Let us now consider an acid-base reaction involving methuylamine that performs as a base and water which acts as an acid in this case, as following below:

Reminder that water can act both as an acid as well as a base. It acts as an acid via donating a proton to yield the OH^- ion which is its conjugate base. It can as well perform as a base via allowing a proton to yield a hydronium ion that is its conjugate acid. Even though, we have demonstrated both the above instances using water as one of the components, the scope of Bronsted-Lowry definition of acids and bases isn't bounded to aqueous results as is the case in Arrhenius definition. The Bronsted-Lowry idea of acids and bases is more common and applies to any kind of solvent.

Therefore, according to this idea the common form of an acid-base reaction can be written as;

Where A₁ - B₁ and A₂ - B₂ are conjugate acid-base pairs.

The acid-base theory was further broadened by Lewis in 1934. He suggested that the acids are the electron-pair acceptors and the bases are the electron-pair donors.

Therefore, according to this thought any molecule or ion that can accommodate an electron pair is an acid. For instance, a proton, H⁺, is a Lewis acid since it can admit an electron pair.

A proton is only one of a huge number of species that might act as a Lewis acid. The electron deficient species these as AlCl₃, BF₃, BCl₃, ZnCl₂, Mg₂⁺ and carbonations are as well Lewis acids. The electron deficient atoms in such species accept the electrons to complete their valence shell octets. Alike any molecule or ion that has an unshared pair of electrons to donate can perform as a base. Therefore, dimethyl ether acts as a Lewis towards boron trichloride which performs like a Lewis acid. This acid-base reaction is symbolized below:

Reminder that the curved arrow shows the movement of a pair of electrons from their source to their destination.

We will agree that the bases are much similar in both the Lewis and the Bronsted-Lowry definitions since a Bronsted-Lowry base must possess a pair of electrons in order to admit a proton. Having recognized a material as an acid or a base according to the above criteria, let us learn how to determine the strength of an acid or a base.

Strengths of Acids and Bases:

It is not possible to determine the strength of an acid or a base in absolute terms.

Therefore, these strengths are always expressed in relative terms. The relative strengths of acids are determined by the extent to which they transfer a proton to a standard base. The standard base which is commonly used for such comparisons is water. Hence, for an acid HA, the proton transfer can be represented by the following equilibrium;

HA + H₂O    ↔     :A^- + H₃O⁺

The equilibrium constant, Keq, for the above equilibrium can be written as,

K_eq = [:A^-][H₃O⁺]/[HA][H₂O]

Wherever the quantities in brackets are the molar concentrations (conveyed as moles dm^-3) of the species at equilibrium. For dilute solutions, the concentration of water is huge and is approximately constant. Therefore, the above expression for equilibrium steady can be rewritten in expressions of a new constant, K_a, termed the acidity constant, as following below:

K_eq [H₂O] = Ka = [:A^-][H₃O⁺]/[HA]

The dissociation of acid HA in solvents other than water can be generalized as,

HA + solvent    ↔     H⁺ - solvent +:A^-

The expression for acidity constant can then be written as follows:

Ka =    [H⁺ - solvent] [: A^-]/[HA]

The acidity constants of different acids have magnitudes ranging from 1014 to 10-50. In order to avoid writing a wide range of powers of 10, K_a is generally expressed in terms of pK_a, where

pK_a.  = - log₁₀K

Taking - log of Eq. and rearranging, we obtain

-log K_a =   -log [H₃O⁺] + log [HA]/[:A^-]

By definition,

-log K_a = , pK_a  and - log [H₃O⁺] = pH

Hence, pK_a = pH + log[HA]/ [:A^-]

This expression relating the pK_a and pH is as well known as Henderson-Hasselbalch equation.  

Therefore, when [HA] = [:A-] then pK_a = pH.

Table: shows the pK_a values for a variety of acids along with their conjugate bases.

Table: Illustrates that the acids which are listed at the top are strong acids. For strong acids these as H₂SO₄, the proton transfer to the base (that is water) is approximately complete and equilibrium lies towards the right.

Therefore, the stronger acids have larger K_a values. Thus, it follows from Eq. 2.4 that the stronger the acid, the smaller the pK_a value. Therefore, as Table demonstrates, the sulponic acids and carboxylic acids are much more acidic as compared to phenol and alcohols.

Remember that the conjugate base of a strong acid will be a weak base and the conjugate base of a weak acid will be a strong base. Likewise, we can simplify for conjugate acids. The stronger the acid, the weaker is its conjugate base and vice versa.

As we familiar that Table lists the pK_a values for protic acids or Bronsted acids only. An analogous Table for the relative acidities of Lewis acids isn't reasonable since in such acids it isn't possible to contain a standard base as reference. But, an approximate order of the strengths of various Lewis acids is as known below:

BF₃            >   AlCl₃                >   FeCl₃         >   SbCl₃        >   ZnCl₂      > HgCl₂

Boron            Aluminium          Ferric         Antimony         Zinc             Mercuric

Trifluoride    chloride                chloride         chloride     chloride           chloride

Lewis acids these as boron trifluoride and aluminium chloride are significant acid catalysts for certain organic reactions.

The Table of pK_a values can be utilized to expect the feasibility of an acid-base reaction. In universal, an acid will transport a proton to the conjugate base of any acid that is below it in the pK_a Table, as well, the larger the difference between the  pK_a values (that is acidities) of the acid and the conjugate acid of  the base, the more favorable will be the proton shift from the acid to the base.

A lot of organic reactions are begin via protonation or deprotonation of a reactant, thus, the pK_a values are as well useful in choosing the appropriate acidic or basic reagents needed for a particular reaction.

Alike to acids, an equilibrium for bases in water can be written a,

:A^- + HOH     ↔     HA  + OH^-

The equilibrium constant for such an equilibrium can be expressed as

K_eq   = [HA] [OH^-]/ [:A^-][HOH]

Where, the quantities in brackets are molar concentrations of the correspond species at equilibrium.

Because the reaction is carried out in aqueous solution, water is acting both as a solvent also as an acid; hence, its concentration can be taken as almost constant. Thus, we can write Eq in terms of the basicity constant, Kb, as

K_eq  =    [H₂O] = K_b = [HA][OH^-]/[:A^-]

The two constants Ka and Ka are related to each other as shown below:

K_aK_b   = [:A^-] [H₃O⁺] / [HA] . [HA][OH^-] /[:A^-] = [H₃O⁺]  [OH^-] = K_w = 10^-14

Where K_w is the self-ionisation constant of water. Hence,

 pK_a + pK_b = 14

Hence, if we know the pK_a of acid HA, the pK_b of the base: A^- can be attained via using the above relation.

In the expressions for K_a and K_b, the concentration of water is normally omitted and therefore, K_a and K_b, have units of moles dm^-3.

The self-ionisation of water can be symbolized as,

H₂O + H₂O   ↔    H₃O⁺   + OH^-

The concentration of the species H₃O⁺ and OH^- in pure water is very low and is equal to 10^-7 moles dm^-2. Hence, the self-ionisation constant, K_w, of water is defined as,

K_w = [H₃O⁺] [OH^-] = 10^-7 x 10^-7 moles² dm^-6 = 10^-14 moles² dm^-6

It is customary to convey the strengths of organic bases not as K_b values but in words of the K_a and pK_a values since it permits a single continuous scale for both acids and bases. As has been stated above the stronger the acid, the weaker will be its conjugate base and vice versa.

In other terms, the stronger the acid, the lower the pK_a, but, the stronger the base, the higher is the pK_a. This is as well marked from Table that whereas the acidity of the acids reduces from top to bottom, the basicity of the conjugate bases amplifies from top to bottom. We can see that NH₂ that comes almost at the bottom of this Table is an extremely strong base (see Table). A comparison of the pK_a values from Table illustrates the subsequent order of the basicities for several of the bases.

-CH₃    > NH₂^-    > RO^-   > OH^- >     RCOO]

As we know that the organic compounds that act as bases can be observed as alkyl derivatives of either water or ammonia; for instance, alcohols (R - O - H), ethers (R - O - R-) and amines RNH₂, R₂NH  and  R₃N. The basic character of such compounds can be featured to atoms these as nitrogen and oxygen which have at least one lone pair of electrons.

Having conversed the strengths of acids and bases, let us now learn the traits influencing the strength of acids and bases. But before proceeding to learn of next section that deals through such factors, answer the subsequent Self Assessment Exercises.

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