Acid Base Theories PDF
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This document details the various acid-base theories, such as Arrhenius, Brønsted-Lowry, and Lewis theories. It covers the properties of acids, bases, and their reactions, along with calculations, examples, and key concepts. The topics encompass the identification of acids and bases, their conjugate pairs, and more complex concepts like acid-base strength and pH.
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Acid Base Theories Baking soda and vinegar have been used as home remedies for generations and today we are only a mouse-click away from claims that baking soda, lemon juice, and apple cider vinegar are miracles cures for everything from cancer to COVID-19. Desp...
Acid Base Theories Baking soda and vinegar have been used as home remedies for generations and today we are only a mouse-click away from claims that baking soda, lemon juice, and apple cider vinegar are miracles cures for everything from cancer to COVID-19. Despite these specious claims, the therapeutic value of controlling acid-base balance is indisputable and is the basis of Food and Drug Administration-approved treatments for constipation, epilepsy, metabolic acidosis, and peptic ulcers. [https://www.ncbi.nlm.nih.gov/pmc/ articles/PMC7544731/] Acid-Base Theories Arrhenius theory An acid is a compound that releases H+ ions in water; and a base is a compound that releases OH– ions in water. Proposed by Savante Arrhenius in 1884. Arrhenius Acids Not all compounds that contain hydrogen are acids. Only a hydrogen that is bonded to a very electronegative element can be released as an ion. Such bonds are highly polar. Arrhenius Bases A base is a compound that releases OH– ions in water Example: NaOH, KOH The base sodium hydroxide (NaOH) – NaOH is an ionic solid. It dissociates into sodium ions in aqueous solution. Some common Arrhenius acid & Base Some Common Acids Some Common Bases Name Formula Name Formula Hydrochloric acid HCl Nitric acid HNO3 Sodium hydroxide NaOH Sulfuric acid H2SO4 Potassium hydroxide KOH Phosphoric acid H3PO4 Calcium hydroxide Ca(OH)2 Ethanoic acid CH3COOH Carbonic acid H2CO3 Magnesium hydroxide Mg(OH)2 A hydrogen atom that can form a hydrogen ion is described as ionizable Nitric acid (HNO3) has one ionizable hydrogen. Nitric acid is classified as a monoprotic acid. Acids that contain two ionizable hydrogens, such as sulfuric acid (H2SO4), are called diprotic acids. Acids that contain three ionizable hydrogens, such as phosphoric acid (H3PO4), are called triprotic acids. Arrhenius Acids Ethanoic acid (CH3COOH), is an example of a molecule that contains both hydrogens that do not ionize and a hydrogen that does ionize. Ethanoic acid is a monoprotic acid. Limitations of Arrhenius Concept Sodium carbonate (Na2CO3) and ammonia (NH3) act as bases when they form aqueous solutions. Neither of these compounds is a hydroxide-containing compound, so neither would be classified as a base by the Arrhenius definition. Brønsted-Lowry Acids and Bases An acid is any molecule or ion that can donate a proton (H+) & a base is any molecule or ion that can accept a proton. Proposed by J.N Bronsted and J.M Lowry in 1923. This theory includes all the acids and bases that Arrhenius defined. It also includes some compounds that Arrhenius did not classify as bases. Superiority over Arrhenius concept Conjugate Acid-Base pairs In an acid-base reaction the acid (HA) gives up its proton (H+) and produces a new base (A–). The new base that is related to the original acid is called a conjugate (meaning related) base. Similarly the original base (B–) after accepting a proton (H+) gives a new acid (HB) which is called a conjugate acid. Examples Conjugate Acids A conjugate acid is the ion or molecule formed when a base gains a hydrogen ion. NH4+ is the conjugate acid of the base NH3. Conjugate Bases A conjugate base is the ion or molecule that remains after an acid loses a hydrogen ion. OH– is the conjugate base of the acid H2O. Conjugate Acids and Bases Conjugate acids are always paired with a base, and conjugate bases are always paired with an acid. A conjugate acid-base pair consists of two ions or molecules related by the loss or gain of one hydrogen ion. Conjugate Acids and Bases In this reaction, hydrogen chloride is the hydrogen-ion donor and is by definition a Brønsted-Lowry acid. Water is the hydrogen-ion acceptor and a Brønsted-Lowry base. The chloride ion is the conjugate base of the acid HCl. The hydronium ion is the conjugate acid of the water base. Conjugate Acids and Bases Strength of Bronsted Acid & Base Some Conjugate Acid-Base Pairs Acid Base HCl Cl– H2SO4 HSO4– H3O+ H2O HSO4– SO42– CH3COOH CH3COO– H2CO3 HCO3− HCO3– CO32– NH4+ NH3 H2O OH– LEWIS ACIDS AND BASES An acid is an electron-pair acceptor & A base is an electron-pair donor. Proposed by G.N Lewis in 1939. Lewis acid base concept is more general than concepts offered by Arrhenius or by Brønsted and Lowry. Identify acid and base NH3 + BF3 → NH3BF3 Ammonia has an unshared pair of electrons to donate. So, it is a base. The boron atom can accept the donated electrons. So, it is an acid. The hydroxide ion can bond to the hydrogen ion because it has an unshared pair of electrons. So, OH− is a Lewis base H+, which accepts the pair of electrons, is a Lewis acid. Superiority of Lewis model of acids and bases All the Bronsted-Lowry acid base reactions are covered by the Lewis model. It is so because the transfer or gain of a proton is accompanied by the loss or donation of an electron-pair in both types of reactions. Many reactions which do not involve transfer of a proton are also covered by the Lewis theory. E.g, NH3 + BF3 → NH3BF3 The Lewis definition is the broadest. It extends to compounds that the Brønsted-Lowry theory does not classify as acids and bases. Acid strength The strength of an acid depends on its ability to transfer its proton (H+) to a base to form its conjugate base. For acid HA, dissolved in water (H2O) For simplicity, Now, we can write the equilibrium reaction as, This equation represents the dissociation of the acid HA into H+ ion and A– ion. So, applying the Law of Mass action to the acid dissociation equilibrium, we can write Here, Ka is called the acid dissociation constant. We know, The strength of an acid is defined as the concentration of H+ ions in its aqueous solution. Therefore, from above equation we can say, , the value of Ka for a particular acid is a measure of its acid strength or acidity Calculation of Relative strength of Weak acids from Ka Exercise Which acid is stronger – HF (Ka = 7.2 × 10-4) or HCN (Ka = 6.2 × 10-10)? Which base is stronger – NH3 (Kb = 1.8 × 10-5) or CH3NH2 (Kb = 4.4 × 10-4)? Can one substance be both an acid and a base? Self-Ionization of Water Although pure water is often considered a nonelectrolyte (nonconductor of electricity), precise measurements do show a very small conduction. This conduction results from self-ionization (or autoionization), a reaction in which two like molecules react to give ions. In the case of water, a proton from one H2O molecule is transferred to another H2O molecule, leaving behind an OH ion and forming a hydronium ion, H3O(aq). pH In chemistry, pH is a scale used to specify the acidity or basicity of an aqueous solution. It is the negative of the base-10 logarithm (log) of the H+ concentration pH = – log [H+] Measurement of pH ✔ Qualitative – Use an indicator to determine if the solution is an acid, a base or a neutral. ✔ Quantitative – Use an indicator or meter to determine the pH of the solution. Indicator An indicator is a chemical that changes colours as the pH of the solution changes. There are many different indicators that are used for a variety of reasons. Many of them are just chemicals that are found in nature. Qualitative Measurement Litmus Litmus is a solid that comes from plant materials. Litmus is typically dissolved in a liquid and then impregnated on paper. Litmus in an acidic solution is red. Litmus in a basic solution is blue. Phenolphthalein Phenolphthalein is a solid that is dissolved in a liquid and typically used as a solution. When phenolphthalein is added to a clear acidic solution, the solution stays clear. When phenolphthalein is added to a clear, basic solution, the solution turns pink or red. Titration Titration is the slow addition of one solution of a known concentration (called a titrant) to a known volume of another solution of unknown concentration until the reaction reaches neutralization, which is often indicated by a color change. In titration the known concentration is used to determine the concentration of an unknown solution. Choice of a suitable indicator The choice of a suitable indicator for a particular acid-base titration depends on the nature of the acid and the base involved in the titration. We may have the titration of : (a) a strong acid with a strong base (b) a weak acid with a strong base (c) a strong acid with a weak base (d) a weak acid with weak base Which indicator is suitable for a given titration, can be found by examining the titration curve of that titration. For Methyl orange - HIn, is red and In- yellow, For phenopthalyn – Hln is colorless and In- is pink pH expression pH Scale The value of Kw found experimentally is 1.0 × 10–14 So, [H+][OH–] = 1.0 × 10–14 [H+][H+] = 1.0 × 10–14 [H+]2 = 1.0 × 10–14 [H+] = 1.0 × 10–7 mol/l Thus the H+ ion and OH– ion concentrations in pure water are both 10– 7 mol/l neutral solution [H+] = [OH–] acidic solution [H+] > [OH–] basic solution [H+] < [OH–] All solutions having pH less than 7 are acidic and those with pH greater than 7 are basic. Relation between pH and pOH We have already stated that pH concept can also be used to express small quantities as [OH–] and Kw. Therefore, pH = – log [H+] pOH = – log [OH–] pKw = – log kw Now, we know Kw = [H+][OH–] log Kw = log [H+] + log [OH–] – log Kw = – log [H+] – log [OH–] pKw = pH + pOH Now, Kw = 1.0 × 10–14 pKw = – log [kw] = – log[1.0 × 10–14] = 14 Since, pKw = pH + pOH; thus pH + pOH = 14 Exercise 1. The hydrogen ion concentration of a fruit juice is 3.3 × 10– 2 M. What is the pH of the juice ? Is it acidic or basic ? 2. If a solution has a pH of 7.41, determine its H+ concentration. 3. Calculate the pH of 0.001 M HCl. 4. If a solution has a pH of 5.50 at 25°C, calculate its [OH–]. 5. Determine the pH of 0.10 M NaOH solution. 6. The pH of a solution of HCl is 2. Find out the amount of acid present in a litre of the solution. 7. If it is requires 75ml of 0.500M NaOH to neutralize 165ml of an HCl solution, what is the concentration of HCl in the solution. Acid Properties Acids are corrosive in nature. They are good conductors of electricity. Their pH values are always less than 7. When reacted with metals, these substances produce hydrogen gas. Example: Mg + H₂SO₄ → MgSO₄ + H₂ Acids are sour in taste. Examples: Sulfuric acid [H2SO4], Hydrochloric acid [HCl], Acetic acid [CH3COOH]. Base Properties They are found to have a soapy texture when touched. These substances release hydroxide ions (OH– ions) when dissolved in water. In their aqueous solutions, bases act as good conductors of electricity. The pH values corresponding to bases are always greater than 7. Bases are bitter-tasting substances which have the ability to turn red litmus paper blue. Examples: NaOH, Mg(OH)2, Ca(OH)2