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Henderson hasselbach equations, common ion effect, strength of acid and bases, le chatlier principle

The document discusses the relative strengths of acids and bases, explaining the concepts of acid dissociation constant (Ka) and base dissociation constant (Kb) according to the law of mass action. It details the common-ion effect and how the addition of an electrolyte can shift equilibrium, and also covers the determination of pH and the Henderson-Hasselbalch equation for buffer solutions. Overall, it presents fundamental principles in acid-base chemistry and their applications.

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Click to edit Master title style 1 RELATIVESTRENGTHS OF ACIDS AND BASES LAW OF MASS ACTION COMMON-ION EFFECTS HENDERSON – HASSELBALCH EQUATION NEHLA P MOULANA COLLEGE OF PHARMACY
Click to edit Master title style 2 RELATIVE STRENGTHS OF ACIDS AND BASES •The stronger acids are those that lose their protons more easily than others. Similarly the stronger bases are those that hold onto protons more strongly than others. •An acid is strong if it completely ionizes in water. In the reaction of Hydrogen chloride with water, for e.g., water acts as a base, accepting the proton from HCI, HCl(aq) + H2O(aq)  Cl- + H3O+ acid base base acid The reverse reaction occurs only to an extremely small extent. Because the reaction goes almost completely to the right, we say that HCI is strong acid.
Click to edit Master title style 3 Relative strength of acids •The strength of an acid depends on its ability to transfer its proton (H+) to a base to form its conjugate base. •When a monoprotic acid (HA) dissolves in water, it transfers its proton to water (a Bronsted base) to form hydronium ion (H3O+) and conjugate base. HA + H2O  H3O+ +A- Acid conjugate base For simplifying our discussion, we take H3O+ = H+ Thus we can write the equilibrium reaction (1) as HA + H2O  H+ +A- ….(2) This equation represents the dissociation of the acid HA into H+ion and A- ion.
Click to edit Master title style 4 Applying the law of mass action to the acid dissociation equilibrium, we can write Ka = [H+][A-] -------------- [HA] (3) [HA] where, Ka is called the acid dissociation constant. It is evident that the concentration of H+ depends on the value of Ka. Therefore the value of Ka for a particular acid is the measure of its acid strength or acidity.
Click to edit Master title style 5 Relative strength of Bases According to the Arrhenius concept, a base is a substance which produces OH ions in aqueous solution. The basic properties of a such a substance are due to these hydroxyl ions. Let us consider a base BOH whose dissociation can be represented as BOH  B+ + OH- Applying the law of mass action to the above equilibrium we can write the equilibrium expression as Kb = [B+][OH-] ----------------- (2) [BOH] Kb is called the base dissociation or base ionisation constant.
Click to edit Master title style 6 •The strength of a base is defined as the concentration OH ions in its aqueous solution at a given temperature. •From the equilibrium expression (2), it is evident that the concentration of OH ions (OH) depends on the value of Kb Therefore, the value of Kb for a certain base is a measure of the base strength.
Click to edit Master title style 7 LAW OF MASS ACTION
Click to edit Master title style 8 LAW OF MASS ACTION Two Norwegian chemists, Guldberg and Waage, In 1867, postulated generalisation called the Law of Mass Action. It states that,”The rate of chemical reaction is proportional to the active masses on the reactants”. By the term “active mass” is meant the molar concentration,. It is expressed by enclosing the formula of the substances in square brackets. E.g., Consider the reaction. CO(g) + 3H2  CH4 + H2O(g) Kc = [CH4] [H2O] ----------------- [CO] [H2]3 Kc = Equilibrium constant.
Click to edit Master title style 9 Consider the general reaction aA + bB--- cC + dD where A, B, C and D denote reactants and products a, b, c and d are coefficients in the balanced chemical equation. For the general reaction, we have Kc = [C]c [D]d Products concentrations ------------ [A]a[B]c Reactants concentrations
Click to edit Master title style 1 0 Hence the law of mass action is a relation that states that the values of the equilibrium-constant expression K, are constant for a particular reaction at a given temperature, whatever equilibrium concentration are substituted.
Click to edit Master title style 11 COMMON-ION EFFECTS •The concentration of a particular ion in an ionic reaction can be increased by the addition of a compound which produces that ion upon dissociation. •The particular ion is thus derived from the compound already in solution and also from the added reagent, hence the name common ion. •In the solution of an electrolyte in water, these exists an equilibrium between the ions and the undissociated molecules to which the law of mass action can be applied. Considering the dissociation of an electrolyte, we have AB ---- A+ + B- and [A] +[B]-/[AB] = k (dissociation constant)
Click to edit Master title style 1 2 • If now another electrolyte or soluble salt is added to the above solution containing a common ion (A+B), it will result in the increase of concentration of the ions A+ or B- and in order that k may remain the same. •The concentration of (AB) must eventually increase. In other words the degree of dissociation of an electrolyte is suppressed by the addition of an another electrolyte (or salt) containing common ion. • Hence, The common ion affect is the shift in an ionic equilibrium caused by the addition of a solute that provides an ion that takes part in the equilibrium. Thus the dissociation of an electrolyte NH4OH is diminished by the addition of NH4Cl which furnishes the common NH4 ion.
Click to edit Master title style 13 THE PH OF A SOLUTION •The pH is defined as the negative of the logarithm of the molar hydrogen ion concentration. Mathematically it may be expressed as. pH = -log [H+] where [H+] is the concentration of hydrogen ions in moles per litre. •A neutral solution, whose hydrogen-ion concentration at 25°C is 1.0 x 10-7M, has a pH of 7.00. •For an acidic solution, the Hydrogen ion concentration is greater than 1.0 x 10-7M, so the pH is less than 7.00. •Similarly, a basic solution has a pH greater than 7.00.
Click to edit Master title style 1 4 Determination of pH of a solution The pH of a solution can be measured accurately by 1. Electrometric method (Potentiometric method) 2. Colorimetric method 1. Potentiometric method The potentiometric method is the more accurate and it is the ultimate standard of Hydrogen ion concentration. In practice, it is applied of means of potentiometric measurements and with a pH meter. The Hydrogen ion concentration of a solution can be determined electrometrically by measuring the potential difference set up between an electrode immersed in the solution and a standard electrode of known potential e.g. standard hydrogen electrode.
Click to edit Master title style 15 Colorimetric Determination of pH •Colorimetric methods are based on the use of reagents which alter colour in accordance with hydrogen ion concentration. These reagents such as litmus (which contains Azolitimic acid, with red undissociated molecules yielding blue anions) are acid base indicators.
Click to edit Master title style 1 6 HENDERSON – HASSELBALCH EQUATION The pH of a buffer solution is related to the pKa of the buffer acid and the log of the ratio of buffering species. This relationship is described by the Henderson-Hasselbalch equation, which can be derived from the ionization constant expression for the weak acid. pH = pKa + log [Conj. Base] ---------------- [Acid] This is an equation relating the pH of a buffer for different concentrations of conjugate acid and base, it is known as the Henderson – Hasselbalch equation.
Click to edit Master title style 17 The weak acid is only slightly dissociated and its dissociation is further depressed by the addition of the salt (N+ A-) which provides A- ions (common ion affect). As a result the equilibrium concentration of the unionised acid is nearly equal to the initial concentration of the acid. The equilibrium concentration [A-] is presumed to be equal to the initial concentration of the salt added since it is completely dissociated. Thus we can write the equation as [H+] = Ka [acid]/[salt] (2) where [acid] is the initial concentration of the added acid and [salt] that of the salt used. Taking negative logs of both sides of the equation (2) we have – log [H+] =- log Ka – log[acid]/[salt] (3)
Click to edit Master title style 1 8 But, – log [H*]=pH and log Ka = pKa, Thus from (3) we have, pH = pKa – log[acid]/[salt] =pKa + log [salt]/[acid] Hence pH = pKa + log[salt]/[acid] This relationship is called the Henderson-Hasselbalch equation or simply Henderson equation.
Click to edit Master title style 1 9 In a similar way, the Henderson equation for a basic buffer can be derived. This can be stated as, POH= pka + log[salt]/[base]