Acids and Bases Acids Acids is substance that
Acids and Bases
Acids • Acids is substance that produce H 3 O+ ions in aqueous solution), HCl (aq) H+ (aq) + Cl- (aq) • H+ can not exist as a separate entity in aqueous solution owing to its strong attraction for the negative pole (the O atom) in H 2 O. The ionization of hydrochloric acid should be written as HCl (aq) + H 2 O H 3 O+ (aq) + Cl- (aq) • The hydrated proton, H 3 O+ is called hydronium ion
Bases is produce OH- ions in aqueous solution. Na. OH (aq) Na+ (aq) + OH- (aq) OH- can accept a proton as follows: OH-(aq) + H+ (aq) H 2 O ( l )
• Other bases are not hydroxides. Instead, they produce OH- ions in water by reacting with water molecules. NH 3 (aq) + H 2 O (ℓ) NH 4+ (aq) + OH- (aq) In a 1. 0 M solution of NH 3 in water, only about 4 molecules of NH 3 out of every 1000 react with water to form NH 4+ and OH-. Thus, when ammonia is dissolved in water, it exists primarily as NH 3 molecules. Nevertheless. Some OH- ions are produced and, therefore, NH 3 is a base.
How Do we Define the Strength of Acids and Bases? Strong acids is one that can react completely or almost completely with water to form H 3 O+ ions. Weak acids produce a much smaller concentration of H 3 O+ ions. CH 3 COOH (aq) + H 2 O (ℓ) CH 3 COO- (aq) + H 3 O+ (aq) Strong base is a base that ionizes completely in aqueous solution. Weak base is a base that ionizes partially in aqueous solution.
Example • Classify each of these species as a BrØnsted acid or base; _ b) NO 2 • A) HBr (aq) _ c) HCO 3 H+ (aq) + Br- (aq) therefore HBr is a BrØnsted acid _ • • • B) NO 2 (aq) + H+ (aq) _ C) HCO 3 (aq) HNO 2 (aq) NO 2 is BrØnsted base 2 - H+ (aq) + CO 3 (aq) _ HCO 3 (aq) + H+ (aq) H 2 CO 3 (aq) Bicarbonate can be act as BrØnsted acid or base is called amphiprotic
What are conjugated Acid-Base Pairs? What happen if acid and base react in absent of water? According to BrØnsted definitions, BrØnsted acids is a proton donor, and BrØnsted Bases is a proton acceptor, and an acid-base reaction is a proton transfer reaction. Furthermore, according to BrØnsted definitions, any pair of molecules or ions that can interconvert by transfer of a proton is called conjugated acid-base pair. When an acid transfers a proton to a base, the acid is converted to its conjugated base. When a base accepts a proton, it is converted to its conjugated acid. Conjugated acid-base pair CH 3 COOH + NH 3 Acetic acid (Acid) Ammonia (Base) CH 3 COO- + NH 4+ Acetate ion (Conjugated base of acetic acid) Ammonium ion (Conjugated acid of ammonia)
Acid Ionization Constant In case of weak acid added to water, HA + H 2 O A - + H 3 O+ The equilibrium constant expression for this ionization is [A-][H 3 O+] k = ________ [HA][H 2 O] As a result of water is the solvent and its concentration changes very little when we added HA to it, we can treat the concentration of water as a constant, so [A-][H 3 O+] ka = ________ [HA] Where ka is the acid ionization constant
• The value of the acid ionization constant for acetic acid is 1. 8 × 10 -5. • p. Ka = - log ka = 4. 75 • Ka tell us how strong an acid. Example Which is the stronger acid: Benzoic acid with a ka of 6. 5 × 10 -5 or hydrocyanic acid with a ka of 4. 9 × 10 -10
Ka and p. Ka for some weak acid Name Ka p. Ka H 3 PO 4 Phosphoric acid 7. 5 × 10 -3 2. 12 HCOOH Formic acid 1. 8 × 10 -4 3. 75 CH 3 CH(OH)COOH Lactic acid 1. 4 × 10 -4 3. 86 CH 3 COOH Acetic acid 1. 8 × 10 -5 4. 75 H 2 CO 3 Carbonic acid 4. 3 × 10 -7 6. 38 H 2 PO 4 - Dihydrogen phosphate 6. 28 × 10 -8 7. 21 H 3 BO 3 Boric acid 7. 3 × 10 -10 9. 41 NH 4+ Ammonium ion 5. 6 × 10 -10 9. 25 HCO 3 - Bicarbonate ion 5. 6 × 10 -11 10. 25 HPO 4 -2 Hydrogen phosphate ion 2. 2 × 10 -13 12. 66 Increasing acid strength Formula
What are the acidic and basic properties of pure water • We have seen that an acids H 3 O+ ions in water, and that a base produce OH- ions in water. • Suppose that we have absolutely pure water; even pure water contains a very small number of H 3 O+ and OH- ions. They are formed by the transfer of a proton from one molecule of water (proton donor) to another (the proton acceptor) • H 2 O + H 2 O Acids Base OH- + H 3 O+ Conjugate base of H 2 O acids of H 2 O [H 3 O+] [OH-] K = _________ [H 2 O] kw = [H 3 O+] [OH-] = 1. 0 × 10 -14 Where kw is the ion product of water
• In pure water, H 3 O+ and OH- form an equal amount, so their concentrations must be equal • [H 3 O+] = 1. 0 × 10 -7 • [OH-] = 1. 0 × 10 -7
p. H is a measure of the acidity or alkalinity of a solution. Aqueous solutions at 25 °C with a p. H less than seven are considered acidic, while those with a p. H greater than seven are considered basic (alkaline). When a p. H level is 7. 0 ; it is defined as 'neutral' at 25 °C because at this p. H the concentration of H 3 O+ equals the concentration of OH− in pure water. Whether a solution is acidic or alkaline, there always both H+ and OH- ions present. It is the predominance of one type of ion over the other that determine the degree of acidity and alkalinity. it is often convenient to define p. H as: Arbitrarily, the p. H is − log 10([H + ]). Therefore, p. H = − log 10[H+]
Examples If one makes a lemonade with a H+ concentration of 0. 0050 moles per litre, its p. H would be: A solution of p. H = 8. 2 will have an [H+] concentration of 10− 8. 2 mol/L, or about 6. 31 × 10− 9 mol/L. A solution with an [H+] concentration of 4. 5 × 10− 4 mol/L will have a p. H value of 3. 35.
p. OH • There is also p. OH, in a sense the opposite of p. H, which measures the concentration of OH− ions, or the basicity.
Measurement • • • by addition of a p. H indicator into the solution under study. An indicator is used to measure the p. H of a substance. Common indicators are phenolphthalein, methyl orange, phenol red, bromothymol blue, bromocresol green and bromocresol purple. by using a p. H meter together with p. H-selective electrodes (p. H glass electrode, hydrogen electrode, quinhydrone electrode, ion sensitive field effect transistor and others). by using p. H paper, indicator paper that turns color corresponding to a p. H on a color key. p. H paper is usually small strips of paper (or a continuous tape that can be torn) that has been soaked in an indicator solution, and is used for approximations. Representative p. H values Substance Hydrochloric acid, 10 M -1. 0 Lead-acid battery 0. 5 Gastric acid 1. 5 – 2. 0 Lemon juice 2. 4 Cola 2. 5 Vinegar 2. 9 Orange or apple juice 3. 5 Tomato Juice 4. 0 Beer 4. 5 Acid Rain <5. 0 Coffee 5. 0 Tea or healthy skin 5. 5 Urine 6. 0 Milk 6. 5 Pure Water 7. 0 Healthy human saliva • The p. H of water gets smaller with higher temperatures. For example, at 50 °C, p. H of water is 6. 55 ± 0. 01. This means that a diluted solution is neutral at 50 °C when its p. H is around 6. 55 and that a p. H of 7. 00 is basic. p. H Blood 6. 5 – 7. 4 7. 34 – 7. 45 Seawater 7. 7 – 8. 3 Hand soap 9. 0 – 10. 0 Household ammonia 11. 5 Bleach 12. 5 Household lye 13. 5
Buffer solutions are solutions that resist change in Hydronium ion and the hydroxide ion concentration (and consequently p. H) upon addition of small amounts of acid or base, or upon dilution. Buffer solutions consist of a weak acid and its conjugate base (more common) or a weak base and its conjugate acid (less common). The resistive action is the result of the equilibrium between the weak acid (HA) and its conjugate base (A−): HA(aq) + H 2 O(l) → H 3 O+(aq) + A−(aq) Any alkali added to the solution is consumed by the Hydronium ions. These ions are mostly regenerated as the equilibrium moves to the right and some of the acid dissociates into Hydronium ions and the conjugate base. Maximum buffering capacity is found when p. H = p. Ka, and buffer range is considered to be at a p. H = p. Ka Buffer solutions are necessary to keep the correct p. H for enzymes in many organisms to work. Many enzymes work only under very precise conditions; if the p. H strays too far out of the margin, the enzymes slow or stop working and can denature, thus permanently disabling its catalytic activity. A buffer of carbonic acid (H 2 CO 3) and bicarbonate (HCO 3−) is present in blood plasma, to maintain a p. H between 7. 35 and 7. 45.
• The most common buffers consist of approximately equal molar amounts of a weak acid and a salt of the weak acid (conjugated base). CH 3 COOH + H 2 O CH 3 COO- + H 3 O+
How Buffers Work H 2 O new HA HA HA A−− Added H 3 O + Tro, Chemistry: A Molecular Approach 19 + H 3 O +
How Buffers Work H 2 O new A− HA HA A−− Added HO− Tro, Chemistry: A Molecular Approach 20 + H 3 O +
How do we calculate the p. H of a buffer? HA + H 2 O A - + H 3 O+ [A-][H 3 O+] Ka = ________ [HA] log Ka = log [H 3 O+ ] + log [A -] _______ [HA] [A -] - log [H 3 O+] = - log [Ka] + log _______ [HA] [A-] p. H = p. Ka + log _______ [HA] Henderson-Hasselbalch Equation
Common Ion Effect HA(aq) + H 2 O(l) A−(aq) + H 3 O+(aq) • adding a salt containing the anion, Na. A, that is the conjugate base of the acid (the common ion) shifts the position of equilibrium to the left • this causes the p. H to be higher than the p. H of the acid solution – lowering the H 3 O+ ion concentration
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