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Wednesday, July 16, 2014

Medicinal Chemistry of SalicylAmide and its Synthesis

 

Synthesis of SalicylAmide

Salicylic acid ==> Salicyl chloride
Salicyl chloride ==> SalicylAmide


Structure Activity Relationship of SalicylAmide 

> Ortho position of hydoxyl (-OH) group is essential for activity
> Removing -OH group decreases its activity
> Substituion of -OH group with other substituents will affect its toxicity and activity
> Acetylation of -OH group will form 2-carbamoylphenyl acetate
> shifting of -OH group to meta or para, the activity will decrease
> substitution of - COOH group with Amide causes decrease or diminish its anti-inflammatory action

Uses

> Antipyretic
>Analgesic

Adrs


> Sedation at high doses
> may cause anemia
> Gastric intolerance


4:04 AM - By Piscean

Tuesday, July 15, 2014

Medicinal Chemistry of Sulfamerazine


Sulfamerazine belongs to antibacterial drugs (bacteriostatic) that contains Sulfonamide as parent structure.
it is used as
  1. bronchitis, 
  2. prostatitis and 
  3. Urinary tract infections
Sulfamerazine as it is belong to Sulfonamide shows following mechanism
It inhibits bacetrial folic acid synthesis by competing PABA with dihydrofolate synthetase

Competitive inhibitor of  dihydrofolate synthetase, this enzyme is used in folate synthesis in bacteria, hence Sulfamerzine interfere with this pathway
 and ultimately decresing the synthesis of bacterial nucleotides and DNA

Synthesis of Sulfamerazine

Synthesis of Sulfamerazine Step 01
(Click to enlarge)

Synthesis of Sulfamerazine Step 02
(Click to enlarge)

Synthesis of Sulfamerazine Step 03
(Click to enlarge)


12:12 PM - By Piscean

Thursday, June 5, 2014

Medicinal Chemistry of Salicylic Acid








11:26 AM - By Piscean

Sunday, June 24, 2012

Resonance and Inductive Effects - Lecture 5



Resonance and Inductive Effects

Functional groups in drug molecules can affect the ionization equilibrium, i.e., can make an acid a stronger or a weaker acid. These electrostatic effects of functional groups are classified as resonance and inductive effects and represent the contribution that resonance stabilization and electronegativity have on ionization.


Inductive Effects
The C-C bond in ethane (CH3-CH3) has no polarity because it involves two equivalent carbon atoms. However, the C-C bond in chloroethane (Cl-CH2-CH3) is polarized by the presence of the electronegative chlorine atom. This polarization is the sum of two effects: first the Cl atom (electronegativity of 3.0) takes some of the electron density of the carbon atom (electronegativity of 2.5) it is attached to, i.e., it polarizes that covalent bond so that the electrons spend more time close to the Cl atom than to the C atom. The carbon attached to chlorine then compensates this withdrawal of electron density by drawing the electrons from the C-C bond closer to itself. This results in the polarization of the C-C bond.



Inductive Effect is the polarization of one bond caused by the polarization of an adjacent bond.

The effect is greatest for adjacent bonds, but can also be felt farther away.

The inductive effect of a substituent in a drug molecule makes it a stronger or weaker acid (relative to the unsubstituted acid), depending on whether the substituent is electron-donating or electron-attracting relative to hydrogen. Polarization of bonds is an effect through space, therefore it decreases in strength the larger the distance from the acidic group.

Functional groups can be classified as electron-withdrawing (-I) or electron-donating (+I) relative to hydrogen. A nitro group (NO2, -I), for example will draw electrons to itself more than a hydrogen atom would if occupying the same position in the drug molecule. Table 2.2 shows functional groups commonly found in drug molecules that can have -I or +I effects.




-I groups increase the acidity of uncharged acids such as acetic acid because they spread the negative charge of the anion. However, -I groups increase the acidity of any acid, no matter what the charge. For example, if the acid has a charge of +1, as in RNH3+Cl-, a -I group destabilizes the positive center by increasing and concentrating the positive charge of the acid. This destabilization is relieved when the proton is lost, therefore the neutral, deprotonated CB form of the acid is favored and the ionization equilibrium shifts towards it. A –I group substitution in RNH3+Cl- makes the parent amine a weaker base.


In general we may say that, groups that withdraw electrons by the inductive effect increase acidity and decrease basicity, while electron-donating groups act the opposite way.


Resonance Effects
Resonance or conjugative effects result from the high mobility of p-electrons and their delocalization in a system of conjugated double bonds. Functional groups that increase the electron density of conjugated systems are called +R (electron-donating), and those that decrease electron density are called –R (electron-withdrawing). An example of resonance effects is found in the higher acidity of carboxylic acids compared to primary alcohols.
Eq.2.18.

Eq.2.19.


The carboxylate anion (RCOO-) is stabilized by resonance not available to the RCH2O- ion or to RCOOH. The RCOO- is stabilized not only by the fact that there are two resonance structures, but also by the fact that the negative charge is spread over two oxygen atoms, rather than concentrated on a single oxygen atom as in RCH2O-. Table 3 shows functional groups commonly found in drug molecules that can have -R or +R effects.



3:45 AM - By Piscean

Strength of Acids and Bases - Lecture 4



Bronsted-Lowry acids vary in their ability to transfer a proton to water upon dissolution, i.e., they vary in their ability to "ionize" in aqueous solutions.

A strong acid is one that can transfer 100% or close to 100% of its acidic hydrogen atoms to water upon dissolution. This process generates many hydronium (H3O) ions in solution, and the equilibrium in equation Eq.2.1. lies far to the right. The actual concentration of unionized species HA present in solution may be so small that cannot be measured, and the resulting Ka (Eq.2.5.) is a very large number.

A weak acid is one that transfers only a small portion of its acidic hydrogen atoms to water upon dissolution. The amount of hydronium ions thus generated are much less than in the case of strong acids. The actual percentage of proton transfer for weak acids depends on the molecular structure, the molecular polarity, and the strength and polarity of individual bonds in the molecule. A weak acid generates a small amount of hydronium ions, and the equilibrium in equation 1 is far to the left, towards the non-ionized species HA. The resulting Ka is small number.
Eq.2.1.

Eq.2.5.


When comparing two acids for their relative strength it is usually the magnitude of the difference (as a power of 10) in their Ka that is most informative. For example,


if acid A1 has a Ka of 5.1 x 105 and acid A2 has a Ka of 3.4 x 102, we can say that acid A1 is approximately 1000 times stronger than acid A2.

This approximation is all we need in many cases. It is also more convenient for our purposes to use the logarithm of Ka rather than Ka itself when comparing acid strengths.

The logarithm of a number is the exponent to which 10 is raised to generate the number:
log10x = x and log10-x = -x

The equilibrium constant for strong acids is a large number, and logarithms of numbers greater than 1 are positive numbers. Strong acids therefore have positive values of logKa.

Weak acids on the other hands, such as many drug molecules, have Ka values in the range between 0 and 1. Logarithm of Ka's in this range are negative numbers.

For example, if acid HA1 donates 1 in 1000 of its acidic hydrogen atoms to water, its Ka will be 1 x 10-3, and the logKa will equal -3. An even weaker acid may donate only one in a million of its acidic hydrogen atoms to water, its Ka will be 1 x 10-6 and its logKa will equal -6.

Scientists have found it more convenient to compare different acids strengths in term of positive numbers, and therefore it is the -logKa that is used to compare different acids. In analogy with pH, which is the -log[H+], the –logKa is called the pKa of the acid. A small pKa means a large Ka (strong acid), just like a small pH means a large [H+].

In the example of the two acids above, the pKa’s will be 3 and 6 respectively, and the acid with the smaller pKa (3) is stronger than the acid with the larger pKa (6).

Many published Tables in the literature list acids and their pKa as a means of comparing their relative strength. Table 2.1 gives some examples of strong acids (large Ka, small pKa) and weak acids (small Ka, large pKa).





Just like there are strong acids and weak acids, there are also strong bases and weak bases. Strong bases are limited to the hydroxides of group IA and IIA in the periodic Table. Weak bases are many of the drug molecules that possess an amine functional group: -NH2. When comparing two bases, the one with a smaller pKb is the stronger of the two. We will not, however, find many Tables published with pKb values to compare base strengths. Instead, scientists have found it more convenient to compare relative acid strengths for all proton-donating species, HA in case of acids and BH+ in case of bases. We will use four common drugs to illustrate:



Fig.2.4.
Weak acids: acetaminophen, aspirin Aspirin
pKa = 3.5 Acetaminophen
pKa = 9.5
Weak bases: amphetamine, diazepam Amphetamine
pKa = 9.8 Diazepam
pKa = 3.3


The comparison of aspirin and acetaminophen is straightforward, as explained before, aspirin with the smaller pKa is stronger than acetaminophen as acid.

In order to compare the two bases we must remember that it is their conjugate acids that we are dealing with, when we use their pKa. The conjugate acid of diazepam, with the small pKa, is stronger than the conjugate acid of amphetamine, which means that diazepam is the weaker base of the two.

In order to make a judicious use of the many pKa tables in the literature one must know the chemical structure of the acidic species listed, and by inspecting the nature of the functional groups present decide whether the pKa refers to an acid or to the conjugate acid of a base. We must know this because for bases, as we saw in the example above, the larger the pKa the stronger the base. This is contrary to acids: the larger the pKa the weaker the acid.

In summary, for calculation purposes, we can view all acid-base reactions in aqueous solutions from the standpoint of the conjugate acid form loosing a proton to form the conjugate base. When we do this we can always use the pKa’s in our calculations and do not need to deal with Kb’s or pKb’s at all. Just remember:

For acids: the stronger the acid, the smaller the pKa
For bases: the stronger the base, the larger the pKa
3:41 AM - By Piscean

Conjugate Acids and Conjugate Bases - Lecture 3



The two acids and two bases involved in a Bronsted-Lowry equilibrium situation can be grouped into two conjugate acid-base pairs.
A conjugate acid-base pair is two species that differ from each other only by one proton.


The two conjugate acid-base pairs in equations Eq.2.1. and Eq.2.2. above are:

Fig.2.1.


The conjugate base of an acid is the species that remains when an acid looses a proton. It is abbreviated CB. The conjugate acid of a base is the species formed when the base accepts a proton. It is abbreviated CA.
3:39 AM - By Piscean

Hydrolysis of Salts - Lecture 2



Hydrolysis of Salts

Many drugs are formulated as salts, a process that involves taking the weak acid or base and reacting it with base or acid, respectively, in order to generate an ionic compound, or salt. When salts are dissolved in water, the ions dissociate completely and associate with water molecules to form solvated anions and cations. One common error is to confuse low solubility with low percent dissociation, but these two processes are totally different.
Barium sulfate, for example, has very low water solubility, but whatever amount does dissolve is 100% ionized into Ba+2 and SO4-2. The same applies to many drugs formulated as salts. They may have varying degrees of water solubility, but whatever amount is dissolved in water is 100% ionized into the component ions. Indeed, the decision to formulate a drug as a salt, usually adding one more step in the manufacturing process, comes from the need to obtain greater solubility in body fluids. The salt is invariably more soluble than the parent compound, although some exceptions will be discussed in a later chapter.

Another caveat to add at this point is that whether the drug is administered as a free acid or base, or as its salt, dissolution in water will bring about the corresponding acid-base equilibrium. The equilibrium constant will be satisfied, regardless of whether we begin with the parent compound or one of its salts. The reaction of potassium acetylsalycilate (aspirin, potassium salt) and water is shown below, abbreviating the acid as HA, the salt as KA.
Eq.2.7.


Once in solution, A- will pick up H+ (from water dissociation) to form HA. We can write this reaction as:
Eq.2.8.



But as soon as H+ is withdrawn from the solution this way, the H2O dissociates some more to replace that H+. This reaction we have seen before:
Eq.2.9.



Adding both these reactions gives us the net reaction for the hydrolysis of A-:
Eq.2.10.



This is the so-called hydrolysis reaction, and its equilibrium constant, commonly designated by Kh (hydrolysis constant), can be written as follows:
Eq.2.11.



We leave out the water from the denominator because as usual its activity stays constant. Multiplying the numerator and denominator by [H+] yields a value for Kh equal to Kw/Ka.

The hydrolysis reaction above will be similar for all salts of drugs that are weak acids and have been reacted with a strong base (i.e. KOH) to form a salt. It also explains why an aqueous solution of the salt of a weak acid is slightly basic, i.e., OH- is generated upon hydrolysis. This last point is not significant in physiological fluids since the latter are commonly buffered, so that they can resist pH changes.

A similar situation arises when the salt we are considering is formed from a parent drug which is a weak base (i.e. epinephrine), and a strong acid (HCl).

Let’s abbreviate epinephrine as RNH2 and its chloride salt as RNH3+Cl- or RNH2.HCl. Upon dissolution of this salt in water there is complete ionization of the salt, as follows:
Eq.2.12.


The Cl- does not affect the water dissociation equilibrium (Eq.2.9), but the RNH3+ can because some of it can combine with OH- to produce RNH2 and H2O by the reaction:
Eq.2.13.


Adding this equation and the water dissociation we get the following net reaction:
Eq.2.14.


This last equation looks like a simple dissociation of a protonic acid, and this is the way we will consider the ionization equilibrium of basic drugs in physiological fluids. In this manner, we can view acids and bases on a similar framework of ionization equilibrium, a perspective which will be convenient for our purposes.

The equilibrium expression for Eq.2.14. has the form:
Eq.2.15.


Multiplying numerator and denominator by [OH-] we get:
Eq.2.16.


Although it would be most proper to call K the hydrolysis constant for the amine RNH2, for the purposes of our discussion we will call it the dissociation constant (Ka) of the cationic form of the amine, or RNH3+.

Equation 2.16 expresses the inverse relationship between the base strength (Kb) of the parent compound, and the acid strength of its protonated, cationic form:


A strong base (large Kb) has a protonated, cationic form which is a weak acid (low Ka)
A weak base (small Kb) has a protonated, cationic form which is a strong acid (high Ka)
3:37 AM - By Piscean

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