Pharmaceutical analysis A summary PDF

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Pharmaceutical analysis A1 ppm (part per million) = 1 μg/mL or 1 mg/L or 1mg/kg 1 ppb (part per billion) = 1 ng/mL or 1 μg/L or 1 μg/kgweight-volume percent (w/v %) = g/100 mL weight percent (w/w %) = g/100 g volume percent (v/v %) = mL/100 mLLecture 1Analytical chemistry : science of obtaining, pro...

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Pharmaceutical analysis A 1 ppm (part per million) = 1 µg/mL or 1 mg/L or 1mg/kg 1 ppb (part per billion) = 1 ng/mL or 1 µg/L or 1 µg/kg weight-volume percent (w/v %) = g/100 mL weight percent (w/w %) = g/100 g volume percent (v/v %) = mL/100 mL

Lecture 1 Analytical chemistry: science of obtaining, processing, and communicating information about the composition and structure of matter (determining what matter is and how much of it exists) Classification of analytic chemistry Method classic wet chemistry -gravimetry: analyte is precipitated by addition of a reagent to form an insoluble solid. The amount of precipitate produced is rinsed, dried and weighed, and the original concentration of analyte is calculated. Exceptionally accurate analysis method, but very laborious to carry out. It is now mostly used to prepare standards for calibration of analytic apparatus -titrimetry: determination of the unknown amount of analyte A, in a solution by measuring the consumption of a reagent, T, that reacts with A (A + T  AT) Equivalentionpoint is amount of reagent necessary to make AT. Reaction must go to completion and fast (all titrant T added must react with analyte) These methods not for small amounts of analyte, not really selective. Instrumental methods (UV-visible absorbance spectrophotometry) Type of information Qualitative: identity, structure Quantitative: amount of analyte in sample Type of system pharmaceutical analysis, forensic and toxicological analysis, food science 1) quality control of medicinal drugs and their dosage forms 2) measurements of bioaccessibility (release of active substance form the drug) 3) determination fo bioavailibility (distribution of active substance and its metabolites in the body) sample is of biological origina and analysis therof is referred to as bioanalysis Quality control European pharmacopoeia reference work for quality control of medicines in Europe, contains official standards (legal and scientific) for control of medicinal products. -all producers of medicines or substances for pharmaceutical use must apply these standards to be permitted to sell their products -Monographs contains: *analytic methods to identify substance or product, and control quality and quantitative strength *quality standards with respect to acceptable levels of impurities (compound having no value)

Analytical process

Laws of equilibrium Law of mass action (Guldberg and Waage) at given temperature, a chemical system reaches a state in which a particular ratio of reactant and product concentration has a constant value.

A,B,C,D: chemical compounds a,b,c,d: indicate how many molecules of a given compound take part in the reaction [A]: concentrations of reacting compounds (mol/L by solution), Pa for gas (unit of pressure) Each of the quantities is actually expressed as the ratio of concentration of the species to its concentration in its standard state -for example: [A] in its standard state [A]stan =1M -therefore: [A] is [A]/(1 M), [B] is [B]/(1 M) = 1 M -for gas: D: Pd is Pd/ (1 bar) Concentrations of pure solids and pure liquids: [A]/[A]stan= 1 Concentrations of solvents (when dissolved analyte concentrations are low): [A]/[A]stan = 1 K is dimensionless, K has no units -when K>1 the numerator is greater than the denominator, equilibrium lies to right (favoured) Stoichiometry: ratios of substances participating in a chemical reaction

Water is an example of a solvent, it has such a high concentration that we can take it as 1, it is not included in the fraction. Le chatelier’s principle / equilibrium law (le chatelier and van ‘t Hoff) When a system at equilibrium is subjected to a change that disturbs it, it will re-adjust itself to partly counteract the effect of the applied change so it can proceed back to equilibrium. Change in concentration from reactant or product

-When dichomate is added (temperature constant), reaction system re-adjusts to the left (reactant and product concentrations shift) until the system’s K value is re-established. -[H+ ] = 5,0 M, [Cr2O7 2-]=0,10 M, [Cr3+] = 0,003 M, [Br- ] = 1,0 M, [BrO3-] = 0,043 M. [Cr2O7 2-] is increased from 0,10 to 0,20 M. In which direction does the system re-adjust itself). The system re-adjust to the left because there is more product. This can be checked with calculating Q (for systems that are not in equilibrium)

Q= K, system in equilibrium Q> K, system re-adjusts to the left Q< K, system re-adjusts to the right Temperature change

Equilibriums for titrimetry Solubiltiy reactions solubility product, Ksp. Example: AgCl (s) ↔ Ag+ (aq) + Cl- (aq)

Ksp is used when an excess of a sparingly soluble salt is immersed in water. both reactants and products are available, equilibirum can be established Salt dissolves until the concentrations of cations and anions have reached

their maximum values  solution is saturated. Under these conditions, the cation and anion concentrations multiplied together equal the Ksp.

Ksp of solution is 1,5x10-9, really small, reaction to the right is not efficient and there will be low concentration of product. Ion effect Concentration of reactants at a sparingly soluble salt change causing the salt to precipitate.

When you add 0,010 M KCl. Cl increased, Ksp is constant so to keep the equilibrium Hg2 needs to decrease. Hg2Cl2 precipitates (Cl from Hg2Cl2 is negligible, because its so small) Acid-base chemistry (bronsted-lowry) -acid dissociation constant, Ka

-base hydrolysis constant, Kb

-acid is a proton doner (gives H+ away) only loses proton if there is a base -base is proton acceptor (picks H+ up) Water can act like an acid or a base

It can also undergo self-ionization: autoprotolysis

Kw is equal to 1,01 x 10-14 (Ksp of water) When a base reacts with an acid, a conjugated acid and a conjugate base are formed. The difference between a conjugate base and the acid is only one H+

Complex formation -stepwise formation constants, K -cumulative formation constants, b

The complexes are stable, this leaves the Kf and β because they are much larger than 1 and so the equilibrium is to right. Equation 1 can go further into Equation 2, these two equations can be added together to form a general one comparison Reduction-oxidation (redox) reactions Ox1 + Red2 ↔ Ox2 + Red1

Lecture 2 Common ion effect

Solubility of Hg2Cl2 (s) decreases, Cl- coming from Hg2Cl2 is negligible Effect of dissolved ions

When KNO3 is added, the number of ions in solution increases. Its so called inert salt (K+ and NO3do not participate in the dissolution reaction). There are more ions in the solution, so the ẟ- or the ẟ + increases so that there is even less attraction between the ions. The tendency of Hg2 2+ particles to interact directly with IO3- particles and precipitate again also decreases. The solubility of Hg2 (IO3) 2 thus increases because the solution is larger [IO3-] and [Hg22+]. Ionic strength

When ionic strength increases: -charge in the ionic atmosphere of ions increases -ions are more shielded -less attraction between cations and anions -reactions producing charged products shift more to the right

The activity coefficient is a measure for how much the behaviour of an ion in a salt solution deviates from the behaviour exhibited by the same ion in an infinitely dilute solution. In other words, gamma is a measure for the influence of the ionic atmosphere on the ion. Gamma for uncharged (neutral) molecules in solution is always equal to unity or 1. In an ideal infinitely diluted solution, the activity coefficient of ion i is equal to 1, the concentration of this ion is (almost) equal to 0. There are therefore no other particles that can influence the ion. In an ideal solution holds KCa = KCT

Not ideal solution:

Right approximation for ionic strength smaller than 0,1 M.

Interpolating

Titration Determination of amount of an analyte A, in a sample by measuring the consumption of a dissolved reagent T, that reacts with A. The quantity of T required for complete reaction of A to AT tells us how much analyte was present. Titration stops when equivalence point is reached. This is the number of mol of T added as titrant is exactly sufficient to react with all the molecules of A (volume titrant added= Veq) Endpoint is addition of titrant is stopped when sudden change in physical property of the sample solution is observed (must be clearly observable, can occur after and before the EQP Titration error= Veqp-Vep

Equilibrium

1) note all important reactions/equilibria 2) write down the full equations for the equilibrium constant

3) write down charge balance 4) write down the mass balance Charge balance: total amount of positive charge in solution must be equal to the total amount of negative charge present Mass balance: quantity of all species in solution containing a particular atom (or groups of atoms) must equal the amount of that atom (or group) delivered to the solution. Statement of the conservation of matter

Lecture Redox titrations

References electrode and indicator electrode, Ecell is difference between those electrodes and is measured after each addition of titrant. The value of the Ecell is dependent on concentrations of the relevant compounds in the solution.

Lecture spectrophotometry Spectroscopy: science that studies the interaction of different kinds of radiation with matter. Spectrophotometry: intensity of radiation (light) is measured and converted to an electronic signal (absorbance and emission). Light as a wave Wavelength : crest-to-crest distance Amplitude A: length of electrical vector at the crest (maximum) of the sine curve Frequency v: number of complete oscillations that the wave makes each second

Speed of light propagation: Vi= v*I, v is constant, if speed changes   changes Electromagnetic spectrum

Speed of light in air is constant c=3.00x10^8 m/s Photon Quantum: physics, the smallest amount or unit of something, especially energy. Amount of energy occur only as multiples, n, of a quantum (n=1, 2, 3) Planck equation: relationship between the wavelength and energy of a photon E = h = hc /  h = Planck’s constant = 6.62  10-34 Js -When light frequency, v, decreases, E decreases -When wavelength, , decreases, E increases Energy states Energy state of atoms, ions, molecules: has to do with the motion of e- around positively charged nucleus (electronic state) -ground state: lowest electronic state -most chemical compounds are in their ground state at room temperature -energy is transferred to electrons by light: energy state changes to a excited state. 1) atoms, ions and molecules can only exist in certain energy state. the energy state of chemical species changes when it loses or gains an amount of energy (photon) which exactly overlaps with the difference in energy between two energy states.

2) relationship between energy difference in a molecule and frequency (wavelength) of the absorbed (or lost) photon: E1 – E0 = h = hc / 

An electronic state is associated with a number of vibrational states. The vibrational states have too with interatomic vibrations. The rotational states have to do with the rotation of molecules. A big transition is a large energy difference is a small wavelength. This is the energy for ultraviolet and visible light of the absorbed photon is mainly used by electronic transitions. Pay attention! An electronic transition is always accompanied by a change in the vibration and rotational states. Every atom, ion and molecule has a unique set of energy states. -qualitative information about molecular structure *chemical groups absorb photons at different frequencies *determining the frequencies of absorbed photons leads to information about chemical groups -Quantitative information about amount of analyte in the solution *determine the number of photons absorbed at a particular wavelength which is specific for an electronic transition.

P: light intensity (energy per unit time per unit cross-section) is given in Watts/m2 Transmittance (fraction transmitted light): T=P/P0 Absorbance: A=log (P0/p)= -log (T) Lambert-Beer Law: A= e*b*c (e= molar absorptivity, c= concentration of the absorbing analyte, b= optical pathlength through sample) Permanent transfer of energy can take place in the interaction between light and matter. In transmission this is not the case! Only a slowdown of the speed of light takes place here. Examples of light-matter interactions without permanent energy transfer are: refraction, dispersion and optical rotation. You can do this with absorption and emission permanent energy transfer takes place.

-Electrons absorb energy from radiation in the UV-Vis regions of the electromagnetic spectrum -Electronic transitions differ, depending on electron type: *we are interested in ,  and n electrons *,  electrons: are bonding electrons *n electrons: lone-pair, non-bonding electrons

Continuous electron motion minimizes repelling force between two positively charged nuclei Single bond: molecular  orbital; double bond: 1x, 1x molecular orbital

The energy transition that consumes the least energy is an n  π * transition. However, this is quantum mechanically prohibited transition, that is, it is not very likely. This is reflected in the low associated molecular weight absorption coefficient ɛ . Probably the π π * feed alley and therefore has a higher ɛ . Other transitions then these two play no role in UV-ZB absorption. Only molecules with double bonds, n, absorb into it UV. The molecules have chromophore groups or simply chromophores.

Electronically excited states: solvent polarity and absorption wavelength. When as a starting point a solution of a chromophore in a non-polar solvent (e.g. hexane) is taken, then it will be added dissolving in a more polar solvent (such as ethanol) for the π  π * to occur a red shift. See figure below. The explanation for this is that the π * orbital has a more polar configuration than the π orbital. The more polar the configuration, the more the orbital is stabilized (energy lowered) by interaction (e.g. H-bridging) with the molecules of the solvent and the lower it gets energy level.

Auxochromes do not themselves absorb light, but modify the ability of chromophores to absorb light: 1) they often cause shifts to longer wavelengths of chromophore absorbance (bathochrome effect) 2) chromophores often exhibit increased absorbance (hyperchrome effect)

Auxochromes do not absorb, but cause shift to longer wavelengths (bathochrome effect) and often an increased absorption (hyperchrome effect)

-First determine the wavelength for the maximum absorption, this is 562 nm -Set the spectrophotometer to 562 nm -Measure the absorbances of solutions of known concentrations at 562 nm as a function of concentration -Fe (II) concentrations in the micromolar range, which is approximately the detection limit for absorbance measurements in most cases. Limitations of absorbance measurements are - Concentrations should not exceed 0.01 M, because above that the molecules of the analyte are too close each other. The charge distribution in each molecule is then changed such that the absorption properties also change. - When the analyte undergoes reactions (or interacts with solvent molecules) products can be formed that have different absorption properties.

Acid-base lecture Polyprotic acids Polyprotic acid had two carboxyl groups loses two protons.

Ka1 is acid with most protons (loss of first proton), Kb1 is base with least protons (gain of first proton)

pH 1) solution with diprotic acid:

Complexometric titrations Detect free metal ions. Complexing agent binds to ion  ethylenediaminetetraacetic acid (EDTA) central ion (+, electron acceptor) + ligand (free electron pair, electron donor) Monodentate ligand: ligand with only one tooth (1 electron pair) Bidentate ligand: more than one lone pair to share Multidentate ligand: form complexes with metal ions that are more stable than complexes formed by similar monodentate ligands....