ACID - BASE EQUILIBRIA

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1 ACID - BASE EQUILIBRIA Mgr. Monika Šrámková Department of medical chemistry and clinical biochemistry, 2 th Medical faculty of Charles Univerzity of Prague and Motol Univerzity Hospital

2 CHEMICAL EQUILIBRIA- INTRODUCTION Chemical equilibria: a) Acid - base b) Redox c) Solubility d) Complexation Law of conservation of mass The mass of reactants is exactly equal to the mass of products AB + CD AD + CB Law of charge conservation The total amount of electric charge entering the reaction is exactly equal to the amount of charge going out of the reaction c z i i 0

3 DYNAMIC EQUILIBRIA At dynamic equilibrium, reactants are converted to products and products are converted to reactants at an equal and constant rate The velocities of reactions are specified by Law of Guldberg Waag Cato Maximilia n Guldberg v 1 = v 2 resp. k 1. c A.c B = k 2. c C.c D Peter Waage

4 EQUILIBRIUM CONSTANT We derive the equilibrium constant from the dynamic equilibrium relation. K k1 k 2 K is concentric equilibrium constant c c If the stoichiometric coefficients in the chemical reaction are not equal to one, chemically reacting: C A *c *c D B k.k + l.l = m.m + n.n

5 EQUILIBRIUM CONSTANT (c M ) m.(c N ) n K = (c K ) k.(c L ) l Equilibrium Concentration Constant is defined as the fraction of the concentrations of reactants promoted to stoichiometric coefficients and the product elevated to stoichiometric coefficients K 1 : in the reaction predominate reactants K 1 : in the reaction predominate products K is close to one: approximately the same amount of reactants and products

6 THE PRINCIPLE OF ACTION AND REACTION- PRINCIPLE OF LE CHATELIERŮV Application of the general principle of action and reaction External action on the chemical system in equilibrium causes the processes that the system tries to eliminate external action External action causes: 1) change of concentration of one component in the reaction 2) change of temperature 3) change of pressure Henry Louis Le Chatelier

7 THEORY OF ACID-BASES (PROTOLYTIC) REACTION The reaction of acids and basis are protolytic reaction Arhenius theory (1884) - acid is a hydrogen-containing compound that, in water produces hydrogen ions (H +) HCl H + + Cl - H 2 SO 4 2 H + + SO 4 2- Base is a hydroxide-containing compund that, in water, produces hydroxide ions (OH - ) NaOH Na + + OH - Ca(OH) 2 Ca OH -

8 THEORY OF ACID-BASES REACTIONS Theory of Brönsted-Lowry (1923) Acid is substance that can donate a protone Base is substance that can accept protone HNO 3 H + + NO 3 - NH 3 + H + NH 4 + HNO 3 + H 2 O H 3 O + + NO 3 - Conjugate Acid-Bases pair is two species, one an acid and one a base, that differ from each other through the loss or gain of a proton

9 THE ION PRUDUCT OF WATER The water is an ampholyte substance that can either lose or accept a proton thus can function as acid or base Autodissosiation (self-ionization) of water: The dissosiation ow water does not influence high concentration of molecular water (55.56mol/l), therefore it can be considered as constant The equation of ion product constant for water is: K v = [H 3 O + ] [OH - ]=[10-7 ][10-7 ]=10-14 The concentrations of both ions are identical at 25 C and 100 kpa

10 ACIDITY AND BASICITY OF SUBSTANCES,VALUE PH According to the equation, H + and OH- are inversely proportional If c (H 3 O + ) > c (OH - ), solution is acidic If c (H 3 O + ) < c (OH - ), solution is basic If c (H 3 O + ) = c (OH - ), solution is neutral Most of physiological solutions have a concentration of hydrogen ions near a neutral point Søren Sørensen (1909) suggested a more practical quantity known as ph ph = log [H 3 O + ] The letter p means negative logarithm of poh= 14 ph

11 DISSOCIATION CONSTANT Describe reactions of acid and bases The constant give an indication of the strenght of acids and bases Its value is given by the tendency of compound to give or to accept a hydrogen ion pk = logk If the value of K is high, it means, that practically 100% of the acid is dissociated. The numeric value of weak acids/bases is low The extent of proton transfer for weak acids is ussualy less than 5%

12 STRENGHT OF ACIDS AND BASES We differ three types of acid and bese acoording to value of dissociation constant: a) Strong - K > 10-2, pk < 2 (HCl, HNO 3,(COOH) 2, NaOH) b) Middle strong K K10-2, pk 2 4 (H 3 PO 4 ) c) Weak K < 10-4, pk > 4 (H 2 CO 3, NH 3, urea ) H 3 O A BH OH KHA K HA B B

13 STRONG AND WEAK ACIDS AND BASES Strong acids are capable of complete dissociation (release of proton), even in a strongly acidic solution Strong acids= high K A and low pk A The weak acids in the acid solution release the proton only partially Bases behave similarly Most of biochemicals compounds behave like weak acids and bases The ph of weak acids and bases must be calculated using a dissociation constant

14 CALCULATION OF PH OF SOLUTIONS OF STRONG MONOTROPIC ACIDS Is acid that supplies one proton per molecule during an acid-base reaction Because the concentration of acid is equal to concentration of H 3 O+, the ph can be obtained from the realtionship: ph= log(h 3 O + ) According to number of hydrogen ion we can classified also diprotic or triprotic acids, than we multiply the concentration by number of hydrogen. ph log 2,3... O H 3

15 CALCULATION PH OF STRONG BASES Process of calculation of ph of strong bases is similar to strong acids poh log OH We obtained value of ph from relationship: ph= 14 poh

16 CALCULATION OF PH OF WEAK ACIDS There are oxonium cations and anions of the acids together with acid and water molecules in aqueous solutions of weak acids and therefore the concentration of oxonium cations is not equal to the concentration of acid 1 ph 2 pk A A If the non-dissociated and dissociated form is of the same concentration, the numerical value of pk is equal to ph The pka value is tabulated, if adds one equivalent of base to one equivalent of acid while experomantal determination of PK, the measured ph is then equal to pk pk A log log K HX ph pk 1 log 1 pk 0

17 DISTRIBUTIVE DIAGRAM It is a graphical illustration of the equilibrium composition of the protolytic system according to ph The diagram is formed by distribution curves Ф = f(ph) Ф is distribution coefficient, unnamed number Ф HA + Ф A = 1

18 DISTRIBUTION DIAGRAM OF ACID- ACETIC ACID, PK=4,755 Ф pkch 3 COOH ph

19 DISTRIBUTION DIAGRAM OF POLYPROTIC ACID-OXALIC ACID pk 1 = 1,25 pk 2 = 4,285 Ф ph

20 CALCULATION OF PH OF WEAK BASES There are hydroxide anions of OH- and cations of the metals together with the base and water molecules In aqueous solutions of weak bases Therefore, the concentration of hydroxide anions is not equal to the concentration of the base ph pk B log MOH pk B log K B

21 HENDERSON-HASSELBACH EQUATION It is derived from the classical dissociation reaction Indicates,if there is more conjugate base than acid in a solution, the ph is higher than pk A If there is more acid than conjugate base, the ph is lower than pak By equation is given relationship between pk A and ph for buffer solution in which is acid-base pair concentration ratio is something other than 1:1 From a practical point of view, the relationship is formulated into ph dependence on the ratio of dissociated and non-dissociated forms (conjugated bases and its acid)

22 HENDERSON-HASSELBACH EQUATION H A K equilibrium constant after crossmultiplication HA H A K HA After a logaritmic calculation and multiplying by -1 of both sides log HA H log K log A Replacing -log [H +] with ph and -log K for pk is obtained ph pk log - A HA

23 BUFFERS Keep ph of a solution from changing very much when either strong acids or strong bases are added to it To handle the addition of both acids and bases, a buffer contains acid (to react with added base) and a base (to react with added acid) Therefore, a combination of an acid-base conjugate pair is used to prepare a buffer Most buffers solution consist of nearly equal concentrations of a weak acid and a salt containing its conjugate base Buffers may also contain a weak base and salt containing its conjugate acid (less common)

24 BUFFERS They are of great importance in maintaining the physiological environment of cells (even in the extracellular space) One of the most effective mechanisms to ensure the appropriate course of biochemical reactions The important systems of buffers are: a) Hydrogencarbonate buffer (HCO 3 - / H2CO 3- ) b) Phosphate buffer(h2po 4 - / HPO 4 2- ) c) System of intracellular proteins The buffering effectis more egffective, if the value of ph of the solution approaches the value of pk. In a simplified way: so it is necessary to add more acid or conjugate base to change value of ph of solution(add or decrease the concentration of proton)

25 THE PH OF AQUEOUS SALT SOLUTIONS Hydrolysis of salts is another example of acidbase equilibria in water solution When acids and bases are mixed, they react with one another and their acidic and basic properties disappear, we say, that they neutralized each other Salts are the products of acid-base neutralization Most of salts disociate to their cation and anion Anion A - derived from weak acid will have reaction like strong base for example We have to decide, if solution of salt is derived from strong or weak acid or base

26 HYDROLYSIS OF SALT We can devided salts into four catgories according to character of their cation and anion a) Type NaCl - salt of strong acid and strong base-does not hydrolyze, the solution is neutral b) Type NH 4 Cl salt of strong acid and weak base hydrolyze to produces an acidic solution c) Type CH 3 COONa salt of weak acid and strong base hydrolyze to produces a basic solution d) Type CH 3 COONH 4 salt of weak acid and weak base hydrolyze to produce a slightly acidic, neutral, or slightly basic solution, depending on the relative weaknesses of the acid and base

27 SOLUBILITY/PRECIPITATION EQUILIBRIA Solubility equilibria exist in systems, where a compound with low solubility is formed A solubility product is a quantitative measure of solubility of a precipitate Ks value defines a limit product of concentrations for which the given precipitate is not forming It specifies an equilibrium of ion concentrations in a saturated solution above the precipitate pks = - log Ks The solubility of precipitation is usually very low ( 10-4 mol l-1)

28 COMPLEXATION EQUILIBRIA They are formed by mixing the metal ion and ligand solutions to form sequentially the coordination compounds The complexes do not arise at one and the same time, but by the sequential occupation of the coordinating spheres of the ion The equilibrium constant for the formation of the complex ion is the formation constant β β = K 1 K 2 K 3.K n The higher the equilibrium constant, the more stable is the coordination complex

29 THANK YOU FOR YOUR ATTENTION