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Medical Chemistry Lecture 1 200 7 (J.S.) Solutions of substances Colligative properties, osmosis Dissociation of electrolytes The equilibria in electrolyte solutions Protolytic reactions Acids and bases , the quantity pH
Liquid dispersions consist of a solvent (dispersion medium, dissolving agent) – usually makes up the greater proportion of the solution and some of the following components (dispersed fraction): dissolved solids (solutes) – molecular compounds – ionic compounds dissociated to ions, colloid particles – macromolecules or micelles particulate materials – aggregates of molecules – precipitated solids, bacteria, cells Analytical dispersions – homogenous, " true“ solutions Colloid dispersions – either colloid sols or colloid solutions Crude dispersions – heterogeneous suspensions and emulsions
Liquid dispersions – three types: < 1nm no no no rapid thermal rapid intensive transparent 1 - 500 nm electron microscope ultrafilters in ultracentrifuge rapid Brownian very slow low max. values opalescent Tyndall effect > 500 nm microscope paper filter spontaneous slow Brownian no no turbid or non-transparent Particle size Particle visibility Filtering capacity Sedimentation Particle movement Diffusion Colligative properties Optical properties ANALYTICAL COLLOID CRUDE Dispersion
Originate in spontaneous dissolving of solutes (either molecular or ionic in size) in solvents,
the term solubility describes the amount of one substance that will dissolve in the solvent,
solubility depends on the polarity of the solute and solvent, on temperature, solubility of gases in liquids also on the pressure of that gas above the solution,
"like dissolves like“ – the most usual solvent of polar compounds is water (weak attraction forces have an important role), nonpolar substances tend to be miscible with nonpolar solvents (e.g., hexane, benzene, diethyl ether, tetrachloromethane).
Analytical dispersions (true solutions)
Concentration of solutions The term dilute or concentrated solutions shows, in a quite relative way, the amount of solute dissolved in a unit volume of solution. The most common types of expressing the concentration of solutions: Amount concentration ( molarity ) , symbol c x , is defined as the amount of the solute (in moles) in one litre of solution, n x / V . Dimension mol l –1 , mmol l –1 , etc. (or M, mM ).
Mass concentration , symbol ρ x (rho)
the mass of the solute (in grams, milligrams, etc.) in one litre of
solution, m x / V ; is used when molarity cannot be applied.
Dimension – g l –1 , mg l –1 , μ g l –1 , etc.
Mass fraction , symbol w x (equal to the percentage by mass) - the ratio of the solute mass to the mass of the solution, m x / ( m x + m solvent ). Comparative relation in mass, without units . Example: Mass fraction w x = 0.05 = 5 % . Such solution contains 5 g of the solute in each 100 g of the solution.
On special occasions, mola l ity is used instead of mola r ity. Molality of a solution, m c , is the concentration expressed as the amount of solute (in moles ) dissolved in 1 kg of solvent , n x / m solvent . Usual dimension – mmol kg H 2 O –1 . For unionized solutes, the colligative properties of a solution are directly proportional to its molality. The value of molality is independent on temperature . It declares the constant ratio between the number of solute and solvent molecules .
What may happen, if a low-molecular weight compound dissolves in water? – Molecules of the compound are dispersed inj water – the solute is a nonelectrolyte . – The compound splits into ions during dissolution - – the formation of ions from a molecular solute (ionization) is only partial , both molecules and ions of the solute are dispersed – the solute is a weak electrolyte . – the compound is ionized or dissociated into ions completely – the solute is a strong electrolyte .
The colligative properties of solutions Colligative properties of solutions: – osmotic pressure, – boiling temperature elevation , – freezing temperature depression , and – solvent vapour pressure lowering are the properties that depend only on the solution molality (on the relative numbers of solute and solvent particles) and not on the identity of particles (their relative masses, shapes, and electric charges).
Osmotic pressure of a solution Water diffuses through the semipermeable membrane from a region of higher water concentration (dilute solution) to one of lower water concentration (more concentrated solution). Pressure Π (pi) is the external pressure exactly sufficient to oppose osmosis and stop it. A semipermeable membrane that allows the free passage of water (solvent) but not the molecules of solutes osmosis Π Π
Measurement of osmotic pressure of solutions by means of osmometers : Calculation of osmotic pressure if the concentration of a solution is known : – membrane osmometers for direct measurements , – osmometers based on cryoscopy - the freezing temperature depression Δ T f is measured and from that value of the solution molality obtained, whi ch is proportional to Π : T f = K f m c . Thermometers have to distinguish changes by 0.001 °C. c – concentration of the solute (in moles) in approximate calculations, for accurate results molality and activity of the solute should be used; i – factor respecting the number of ions formed by dissociation of the solute; R – ideal gas constant equal to 8.314 J /(K mol); T – temperature (in kelvins). = i c R T (in kilopascals)
The value of i for non - electrolytes equals 1, i = 1 for strong electrolytes is a whole number greater than 2, i > 2 α c – degree of dissociation at the concentration c N – number of ions resulting from dissociation of the formal unit Examples: NaCl Na + + Cl – i = 2 Na 2 SO 4 2 Na + + SO 4 2– i = 3 MgCl 2 Mg 2+ + 2 Cl – i = 3 Na 3 PO 4 3 Na + + PO 4 3– i = 4 Factor i in the equation = i c RT : For weak electrolytes i = 1 + α c c ( N – 1 )
All various solutions that have the same osmotic pressure, because of the same osmolality , are isotonic to each other . Hypertonic solutions are solutions at higher osmolality than those to which they are compared; similarly, hypotonic solutions are those with lower osmotic pressure. Solutions isotonic with blood plasma have osmotic pressure about 765 kPa (osmolality about 290 mmol/kg H 2 O ). In order to prevent possible injury to blood cells by osmosis, fluids for intravenous use are usually prepared at approx. isotonic concentration. Sodium chloride solution isotonic with blood plasma (called inaccurately " physiological saline solution“) contains 154 mmol NaCl per litre (154 mmol/l Na + and 154 mmol/l Cl – ), that is 9 g NaCl / l (0.9% sodium chloride). Glucose solution isotonic with blood plasma contains 308 mmol glucose per litre , that is 55 g / l (5% glucose solution).
Cytolysis A cell in an isotonic fluid H 2 O H 2 O H 2 O Osmosis in a hypertonic fluid – a cell shrinks H 2 O H 2 O H 2 O A cell swells in a hypotonic fluid
Blood plasma osmolality is measured by means of osmometers. Blood plasma osmolality 280 – 295 mmol kg H 2 O –1 Hypoosmolality (up to 230 mmol/kg) – deficit in Na + or hyperhydration Hyperosmolality (up to 400 mmol/kg) – retention of Na + , dehydration, hyperglycaemia, uremic syndrome, unusual compounds (e.g. ethanol, ethylene glycol, acetone). It is under the strict hormonal control (aldosterone, vasopressin, atrial natriuretic peptides). Plasma osmolality (mmol kg H 2 O -1 ) ≈ 2 [ Na + ] + [ glucose ] + [ urea] (mmol/l) or ≈ 1.86 [Na + ] + [glucose] + [urea] + 9 (mmol/l) The marked difference between the measured and the roughly estimated value is the sign of the " osmolar gap“ that is usually caused by high concentration of other unionized compou n d s (ethanol, acetone, etc.). In spite of the known value of osmolality gained by measurement, it is useful to calculate the rough estimate of plasma osmolality from the values of major plasma solutes:
Osmotic pressure of molecular colloid solutions (oncotic pressure) is very small when compared to that of true solutions; because of large size of molecules, colloid solution of high- molecular compounds cannot reach high osmolality values . Oncotic pressure of blood plasma proteins represents less than 0.5 % of the oncotic pressure of blood plasma . In spite of this low value, oncotic pressure is extremely important for shifts of water between blood plasma and interstitial fluid in blood capillaries.
When certain low-molecular weight substance dissolves in water and – molecules of the compound are dispersed, the solute is a nonelectrolyte (the solution is a nonconductor); – ions of the compound exist the solution that is a conductor of electricity, the solute is an electrolyte . – if the formation of ions from a molecular solute (ionization) is only partial , the solute is a weak electrolyte , both molecules and ions of the solute are present, – if the compound splits into ions completely due to dissociation or ionization, the solute is a strong electrolyte .
Electrolytes are solutions of compounds that are split into ions due to interaction with a polar solvent. Ionic compounds dissociate completely, polar molecular compounds are ionized completely . I ons are surrounded by a certain number of water molecules (hydrated). Concentration of particles is higher ( in strong electrolytes at least two times, if only two ions are formed) – remember the colligative properties and factor i ! Weak electrolytes: AB ( s ) A + ( aq ) + B – ( aq ) + AB ( aq ) H 2 O Weak electrolytes are ionized to only a slight extent, the concentration of ions are relatively low; most solute molecules do not split into ions. AB ( s ) A + ( aq ) + B – ( aq ) H 2 O Strong electrolytes:
strong acids strong hydroxides most soluble salts Strong electrolytes Weak electrolytes weak acids weak bases ( a few salts) Strong acids: H 2 SO 4 , HNO 3 , HCl , HBr , HI HClO 3 , HClO 4 , alkyl sulfates, alkanesulfonic acids Strong hydroxides: NaOH , KOH , Ca ( OH ) 2 , Sr(OH) 2 , Ba(OH) 2 , tetraalkylammonium hydroxides Water and exceptions among salts, e.g. calcium citrate, ZnCl 2 , HgCl 2 Weak acid: all not named among the strong (nearly all organic acids included) Weak bases: ammonia all other nitrogenous bases
Strong electrolytes Concentrations of ions in solutions of strong electrolytes is always higher than the concentration of the compound. For example: c (Na 2 SO 4 ) = 0.1 mol / l In this solution c (Na + ) = 0.2 mol / l and c (SO 4 2– ) = 0.1 mol / l, i.e. c (Na + +Cl – ) = 0.3 mol / l Due to the mutual electrostatic interactions at higher concentrations (exceeding 10 –4 mol / l), there are certain differences in the behaviour (in colligative properties) of strong electrolyte solutions. The solutions seem to be more diluted than the real concentration proved by chemical analysis. Real colligative properties correspond more with the quantity called activity of ions than with the " analytical“ concentration of ions.
a i activity of the ion i i activity coefficient of ion i at c i c i concentration of ion i ( 1 ) Only for c i < 10 –4 mol/l y ic 1 and a i = c i Activity of ions is a quantity representing the concentration of ions corrected for interionic interactions . Activity coefficients take values up to 1 (= no difference between activity and concentration). In most cases that will be dealt with in this course, the difference between c i and a i will be neglected (what is fully true for concentrations lower than 10 –4 mol/l ) . a i = i c c i
Examples of the activity coefficient values The mean ion activity a ± = y ± c i The values of activity coefficients y ± : HCl 0.97 0.92 0.78 NaCl 0.95 0.83 0.61 H 2 SO 4 0.81 0.61 < 0.50 Cations and anions of the strong electrolyte type MX 2 c = 0.01 mol/l c = 0.1 mol/l c = 1 mol/l n (MX 2 )/ n (H 2 O) 1 : 5550 1 : 555 1 : 56
Ionic strength of solutions I is the function of ion concentration and electric charge and id defined by the relation c i concentration of ion i , z i electric charge of ion i The value of an activity coefficient depends on the concentration and electric charge of all ionic species in the solution, e.g. for concentrations up to 10 –2 mol/l – log y ic = A z i 2 , where is a new quantity I named. This quantity cannot be measured, it can be only calculated: Not all ions exhibit the same effect in a solution, polyvalent ions play a greater role than monovalent ions. I = ‒ 1 2 c 1 z 1 2 + c 2 z 2 2 + … c n z n 2 = 1 2 ‒ c i z i 2 i
In various types of salts and other strong electrolytes, the relations between concentration of the compound, concentration of ions, and ionic strength are different. Examples: Type of salt c salt c particles Ionic strength I Na + Cl – c salt 2 c salt c salt Ca 2 + Cl 2 – c salt 3 c salt 3 c salt Zn 2+ SO 4 2– c salt 2 c salt 4 c salt Fe 3+ Cl 3 – c salt 4 c salt 6 c salt
Various types of equilibria in electrolyte solutions Four types of equilibria may occur in aqueous electrolyte solutions: 1 Protolytic equilibria (acid-base equilibria) deal with the exchange of protons (hydronium ions) between acids and bases 2 Complex-forming equilibria exist between donors and acceptors of valence shell electron pairs 3 Dissolution equilibria express relations of solids to ions and polar solvents in saturated solutions 4 Oxidation-reduction equilibria deal with the exchange of electrons in redox reactions
Acids and bases pH values of acids and bases solutions Protolytic reactions
Acids and bases according to the Br ønsted concept : Acids are proton donors . Any molecule or ion that can lose a proton , if a base is present to accept it, is an acid. Bases are proton acceptors . Any molecule or ion that have an unshared electron pair able to bind a proton released by an acid by coordinate bond is a base. HA + H 2 O A – + H 3 O + acid water as a base B l + H 2 O B H + + OH – base water as an acid
Conjugate pairs Each acid after releasing a proton becomes a base that is called a conjugate base of the primary acid. Similarly, each base gives its conjugate acid by accepting a proton. Therefore, in any reaction of an acid with a base, two conjugate pairs must take part: The first conjugate pair: HA -> H + + A – The second pair: B + H + -> BH + HA + B A – + BH + conjugate pair conjugate pair
Notice that can help you in understanding the following matter and might be appreciated in the study of biochemistry, physiology, and clinical applications of protolytic reactions: Strong hydroxides dissociate to the strong base OH – (conjugate base of water) and cations Na + , K + , Ca 2+ , etc. Those cations are not acids, cannot release H + , and do take part in protolytic reactions. They are called "strong" or spectator cations . Quite generally, acids dissociate to give H + and conjugate bases. Weak acids give H + and strong conjugate bases that exhibit a great affinity to H + ; therefore, only a small part of molecules dissociates in aqueous solution. Conjugate bases of strong acids , anions Cl – , SO 4 2– , NO 3 – , etc., are, on the contrary, exceedingly weak conjugate bases in the body (at physiological values of pH) . They do not notice H + (they do not take part in protolytic reactions) and are named "strong" or spectator anions . Weak bases accept protons only to a small extent and give so strong conjugate acids .
Ionization of water Water molecules are polar and ionize to a very slight extent. Pure water is a very weak electrolyte . One molecule of water gains a proton forming thus a hydronium ion while another molecule forms a hydroxide anion . Water molecules can act as either an acid or a bases; such species of particles are named amphoteric (or amphiprotic). 2 H 2 O H 3 O + + OH – The symbol H + is commonly used in describing protolytic reactions though there are no free protons in water or aqueous solutions; H + must be always taken as a simplified notation of hydronium ions H 3 O + (as well as of more hydrated H + .nH 2 O ions).
Ionic product of water K w From the quantitative point of view, the ionization of water is described by the equilibrium ionization constant K c : Because the high value of unionized molecules of water [H 2 O] (55.5 mol/l) changes only slightly in real solutions, instead of K c the constant ionic product of water K w is used: In any aqueous solutions, concentrations [ H + ] and [ OH – ] can achieve only those values that are in agreement with the K w value 1.00 10 –14 . K c [H + ] [OH – ] [H 2 O] K c [H 2 O] = K w = [ H + ] [ OH – ] = 1.00 10 –14 (25 °C)
In pure water , the concentrations [H + ] and [OH – ] are the same and the concentration of each of this ions is equal to 1.00 10 –7 mol/l (at 25 °C) . All aqueous solution in which [H + ] equals [OH–] are neutral . In acidic solutions [H+] is greater than [OH – ], in alkaline solutions (basic solutions) [OH – ] > [H + ]. Based upon Le Chatelier´s principle, adding H + to a neutral aqueous solution will decrease [OH – ] and similarly, adding OH – will decrease [H + ]. A more convenient way for expressing [H + ] and [OH – ] in mol/l (small numbers, scientific notation) are the expressions pH and pOH defined as the negative logarithms of [H + ] and [OH – ]: pH = – log [ H + ] and pOH = – log [ OH – ]
The values of pH and pOH for dilute aqueous solutions fall between 0 and 14 . The same logarithmic notation can be also used for K w : K w = [H + ] [OH – ] = 1.00 10 –14 p K w = pH + pOH = 14 so that pH = 14 – pOH The pH and pOH scales are logarithmic, not linear, scales ! A change of pH by 1 represents the 10 times higher or lower concentration of [H + ]; any two-fold increase in [H+] results in the decrease of pH by 0.3, because the logarithm of 2 equals 0.30.
Dissociation of a strong acid: HA + H 2 O H 3 O + + A –
In solutions of monoprotic strong acids , the concentration [H + ] is equal to the total strong acid concentration c HA :
[ H + ] = c HA and pH = – log [ H + ] = – log c HA
Dissociation of a strong hydroxide: MeOH( aq ) Me + + OH –
pH of strong acids and strong hydroxides solutions Strong acid and strong hydroxides are strong electrolytes that are fully dissociated in aqueous solutions. (In diluted solutions, the difference between the concentration and the activity of ions can be neglected.) Let us remind the list of a few strong acids and hydroxides!
The titration curve of a monoprotic strong acid ( c HA = 0.1 mol/l) Linear decrease in [H + ] in the course of titration results in a logarithmic curve representing the increase of pH values. To reach the pH value greater by 1 than the initial, 90 % of the amount of an strong acid have to be neutralized. 12 10 8 6 4 2 0 pH pH 7.00 (excess NaOH) n ( OH – ) / n ( acid ) 0 0.2 0.4 0.6 0.8 1.0
Dissociation of weak electrolytes, ( of weak acids and weak bases ) A weak monoprotic acid A weak base Equilibrium constants of ionization Acid and base ionization constants Weak electrolytes are ionized to only a slight extent, the concentration of ions are relatively low; most solute molecules do not split into ions. HA + H 2 O H 3 O + + A – B + H 2 O BH + + OH – K c [H 3 O + ] [A – ] [H 2 O] [HA] K c [BH + ] [OH – ] [H 2 O] [B] The extent to which the weak electrolytes dissociates is expressed by either ionization constants or degree of ionization. K A [ H + ] [ A – ] [ HA ] K c [H 2 O] = K B [ BH + ] [ OH – ] [ B ]
Displacement of a weak acid The consequence of the Le Chatelier´s principle is that acidifying of the solution of a weak acid by adding a stronger acid will suppress the dissociation of the weak acid . If a weak acid is volatile (e.g. HCN, H 2 CO 3 ), it can leak completely from the solution. Displacement of weak bases Like acids, the dissociation of which is kept down by addition of strong acids, alkalization of solutions suppresses ionization of weak bases . Weak acids and weak bases can be also displaced from the solutions of their salts.
p K = – log K K A and K B are acid and base ionization constants. Instead of them, the values of p K A and p K B are used: The lower the value of p K A , the stronger is the weak acid . " Moderately strong " weak acids and bases have p K value from 1 to 3, weak acids and bases from 4 to 8, very weak acids and bases higher than 8. The relation between K B and K A of weak bases For any particular weak base in aqueous solution it holds that p K B + p K A = 14 = p K w K B K A = 1. 10 -14 = K w
p K A of weak acids Oxalic (COOH) 2 1.25 . 4.29 - H 3 PO 4 2.16 7.20 12.29 HNO 2 3.39 - - Ascorbic 4.17 11.57 - Acetic CH 3 COOH 4 . 76 - - H 2 CO 3 6 . 35 10.3 - H 2 S 7 . 07 12.2 - H 3 BO 3 9 . 24 12. 7 - p K B and p K A of weak bases The weaker is the acid, the stronger conjugate base is its anion. The weaker is the base, the stronger conjugate acid is its cation. Weak acid p K A 1 p K A 2 p K A 3 Weak bases p K B p K A conjugate acid Guanidine 1.50 12.50 Methylamine 3.36 10.64 Ammonia 4.75 9.25 Imida zole 6.90 7.10 P yridine 8.82 5.18 Aniline 9.38 4.64. Coffei ne 13.40 0.60
The second parameter that expresses the dissociation (ionization) of weak acids and weak bases is the degree of ionization α c (the percentage of ionization). It is found from the concentration of H + (in acid solutions, of OH – in solutions of bases) and the total concentration of the acid or base c total that is equal to the sum of both ionized and non-ionized molecules [H + ]+[HA] or [OH – ]+[BH]. For a monoprotic weak acid α c = [ H + ] c total The index specifies the total concentration, the value of α varies with the c total . The more dilute the solution, the larger is the degree α c . Because [H + ] = α c c tota l , then K A = c total α 2 1 – α If the degree of ionization is low ( α c < 0.10, i.e. 10 %), the relation is simpler: K c total α c 2 and α c ( K / c total ) ½ .
Calculation of pH values in solutions of weak acids and bases Solutions of weak monoprotic acids The total acid concentration c total = [H + ] + [HA] . Because [H + ] = [A – ] (molecules of acid gives the same number of both), [H + ] [A – ] = [H + ] 2 . Now an approximation is induced: [HA], the concentration of undissociated molecules ( c total – [H + ]), is close to c total , when the dissociation degree α c is less than 0.10 (the minute difference of the denominator can be neglected). Then , from which and in logarithmic form log [ H + ] = ½ log K A + ½ log c total The pH value of a weak acid solution HA + H 2 O H 3 O + + A – K A [ H + ] [ A – ] [ HA ] K A [ H + ] 2 c total [ H + ] = pH = ½ p K A – ½ log c total
pOH = ½ p K B – ½ log c B Solutions of weak bases The total base concentration c total = [BH + ] + [B] . Because [BH + ] = [OH – ] (molecules of base gives the same number of both), [BH + ] [OH – ] = [OH – ] 2 . The approximation as for acids: [B], the concentration of undissociated base ( c total – [OH – ]), is close to c total , when the dissociation degree α c is less than 0.10 (the minute difference of the denominator can be neglected). Then , from which and in logarithmic form log [ OH – ] = ½ log K B + ½ log c total The pH value of a weak base solution B + H 2 O BH + + OH – K B [ BH + ] [ OH – ] [ B ] K B [ OH – ] 2 c total [ OH – ] = pH = 14 – ½ p K B + ½ log c B