Acid-Base Theories In defining what is considered to be an acid and what is considered to be a base, three theories have been proposed: Chapter 15 Acids and Bases Arrhenius acid-base theory Brønsted-Lowry acid-base theory Lewis acid-base theory Dr. Peter Warburton [email protected] http://www.chem.mun.ca/zcourses/1011.php We will see that each subsequent theory builds upon what was stated in the previous theory. 2 The Arrhenius Theory Arrhenius Theory – Acid Strength In the Arrhenius theory of acids, an acid dissolved in water increases the concentration of hydronium ions H3O+ in the solution: Arrhenius acid HA: HA (aq) + H2O (l) ' H3O+ (aq) + A- (aq) In this reaction we see all Arrhenius acids contain protons (H+) that are donated to water In the Arrhenius theory of acids, a strong acid COMPLETELY reacts with water, so there is no HA left at the end of the reaction: HA (aq) + H2O (l) 3 →H O 3 + (aq) + A- (aq) 4 1 Arrhenius Theory – Acid Strength The Arrhenius Theory In the Arrhenius theory of acids, a weak acid reacts with water until an equilibrium is reached where HA is still present in the equilibrium mixture: HA (aq) + H2O (l) ⇌HO 3 + In the Arrhenius theory of bases, a base dissolved in water increases the concentration of hydroxide ions OH- in the solution: Arrhenius base M(OH)x: M(OH)x (aq) (aq) + A- (aq) ' Mx+ (aq) + x OH- (aq) In this reaction we see all Arrhenius bases contain hydroxide (OH-) 5 Arrhenius Theory – Base Strength Arrhenius Theory – Acid Strength In the Arrhenius theory of acids, a strong base COMPLETELY dissociates in water, so there is no M(OH)x left at the end of the reaction: M(OH)x (aq) 6 → Mx+ (aq) + x OH- (aq) 7 In the Arrhenius theory of bases, a weak base only partially dissociates in water until an equilibrium is reached where M(OH)x is still present in the equilibrium mixture: M(OH)x (aq) ' Mx+ (aq) + x OH- (aq) 8 2 Why do we need to improve on Arrhenius theory? Common strong acids and bases The Arrhenius theory has a drawback! Certain compounds that DO NOT contain hydroxide can still increase the hydroxide concentration when placed in water. Arrhenius theory does not explain this! 9 10 Proton transfer reactions The Brønsted-Lowry Theory The Brønsted-Lowry Theory: an acid Brønsted-Lowry Theory: an acid is any substance that donates protons (H+) while a base is any substance that can accept protons. This means that Brønsted-Lowry acid-base reactions are proton transfer reactions. 11 Pairs of compounds are related to each other through Brønsted-Lowry acid-base reactions. These are conjugate acid-base pairs. 12 3 Proton transfer reactions Water in BL acid-base reactions When a Brønsted-Lowry acid is placed in water, it donates a proton to the water (which acts as a base) and establishes an acid-base equilibrium. Generally, an acid HA has a conjugate base A- (an H+ has transferred away from the acid). Conversely, a base B has a conjugate acid BH+ (an H+ has transferred toward the base). 13 Water in BL acid-base reactions 14 Water in BL acid-base reactions When a Brønsted-Lowry base is placed in water, it accepts a proton from water (which acts as an acid) and establishes an acid-base equilibrium. In the reverse reaction of the equilibrium, the acid H3O+ donates a proton to the base A- to give back water and HA. 15 16 4 Water in BL acid-base reactions Brønsted-Lowry Bases To accept a proton (to act as a B-L base) requires a molecule to have an unshared pair of electrons which can then be used to create a bond to the H+. In the reverse reaction of the equilibrium, the acid BH+ donates a proton to the base OH- to give back water and B. All Brønsted-Lowry bases have at least one lone pair of electrons. 17 18 Amphiprotic substances Brønsted-Lowry Bases In the previous reactions we’ve seen NH3 has a lone pair of electrons and can act as a B-L base. Also, water has two lone pairs of electrons, and can act as a B-L base. 19 Some substances, like water, have protons that can be donated (BL acid), and lone pairs of electrons that can accept protons (BL base). This is why it can act like an acid OR a base DEPENDING on the other species present. Such substances are said to be amphiprotic. 20 5 Problem Problem Write a balanced equation for the dissociation of each of the following Brønsted-Lowry acids in water: What is the conjugate acid of each of the following Brønsted- Lowry bases? a) H2SO4 b) HSO4c) H3O+ d) NH4+ a) HCO3b) CO32c) OHd) H2PO421 Why do we need to improve on Brønsted-Lowry theory? 22 Lewis Acids and Bases A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. There are many reactions that behave VERY MUCH LIKE proton transfer reactions that DO NOT involve protons! These definitions are more general than the Brønsted-Lowry definitions because protons DO NOT need to be involved in Lewis acid-base reactions. 23 24 6 Metal ions as Lewis acids In general LA + :LB Æ LA-LB BF3 is a Lewis acid where the B atom can accept an electron pair. The N of the NH3 has a lone pair that can be donated, making NH3 a Lewis base. Many metal ions have the ability to act as Lewis acids. The ions are willing to accept electron pairs from LIGANDS (which act as Lewis bases) because this often stabilizes the ion in solution. The result is often called a complex ion. 25 Problem 26 Comparing theories Since Arrhenius acids must contain protons, then ALL Arrhenius acids ARE ALSO Brønsted-Lowry acids. We’ve already seen that NOT ALL Brønsted-Lowry bases are Arrhenius bases. 27 28 7 Comparing theories BL acid and base strength There are Lewis acids (like metal ions) that ARE NOT Brønsted-Lowry acids. Brønsted-Lowry acid-base equilibria are competitions! ALL Brønsted-Lowry bases must all have at least one lone pair of electrons, so The equilibrium is the result of a tug-of war between the two bases in the system as they fight for protons given away by the two acids. ALL Brønsted-Lowry bases MUST ALSO BE Lewis bases 29 30 Strong BL acids in water Acid Strength and Base Strength The acid that is “better at donating protons” OR the base that is “better at accepting protons” will be found in lesser amounts at equilibrium compared to the other acid (or base). 31 A strong acid (HA) is one that almost completely dissociates in water (which acts as a base). The conjugate base A- will be a very weak base. ← HA + H 2 O ⎯⎯→ A − + H 3O + 32 8 Strong BL acids in water Weak BL acids in water ← HA + H 2 O ⎯⎯→ A − + H 3O + At equilibrium, there will be very little to no HA present in the system, and the concentration of A- will essentially be the same as the initial concentration of HA. A weak acid (HA) is one that partially dissociates in water (which acts as a base). The acid is not as good at donating protons to the water. The conjugate base (A-) will be a weak base. Overall HA + H 2 O → A − + H 3O + ← 33 34 Weak BL acids in water − + HA + H 2 O → A + H O 3 ← At equilibrium, there will be some Aand H3O+ present in the system. However, the concentration of HA will still be significant at equilibrium. 35 36 9 Hydrated Protons and Hydronium Ions Notice that the strongest acids have the weakest conjugate bases, and the strongest bases have the weakest conjugate acids! The ultimate proton-donor is a proton itself! In water there is no such thing as H+. H + + H 2O ⎯ ⎯→ H3O+ Often more than one water molecule will crowd around the proton to give hydrates with the formula H(H2O)n+ where n is 1 to 4. 37 38 Dissociation of Water Hydrated Protons and Hydronium Ions It is possible for one water molecule to act as an acid while another water molecule acts as a base at the same time. This leads to the selfionization of water equilibrium: H2O (l) + H2O (l) ' H3O+ (aq) + OH- (aq) The equilibrium constant for this reaction is called the ion-product constant for water, Kw. Kw = [H3O+][OH-] 39 40 10 At 25 °C Acidic [H3O+] > 1.0 x 10-7 M or [OH-] < 1.0 x 10-7 M Basic [OH-] > 1.0 x 10-7 M or [H3O+] < 1.0 x 10-7 M Neutral [H3O+] = [OH-] = 1.0 x 10-7 M At 25 °C, Kw = 1.0 x 10-14 so [H3O+] = [OH-] = 1.0 x 10-7 mol/L Relatively few water molecules are dissociated at equilibrium at room temperature! We will always assume that [H3O+] [OH-] = 1.0 x 10-14 at 25 °C. 41 42 At 25 °C At 25 °C We also find, since [H3O+] [OH-] = 1.0 x 10-14 = Kw then [H3O+] = 1.0 x 10-14 / [OH-] and [OH-] = 1.0 x 10-14 / [H3O+] 43 44 11 Problem Problem The concentration of OH- in a sample of seawater is 5.0 x 10-6 mol/L. Calculate the concentration of H3O+ ions, and classify the solution as acidic, neutral, or basic. At 50 °C the value of Kw is 5.5 x 10-14. What are the [H3O+] and [OH-] in a neutral solution at 50 °C? 45 The pH Scale 46 pH and acidity [H3O+] in water can range from very small (strongly basic) to very large (strongly acidic) it is sometimes easier to use a negative logarithmic (power of 10) scale to express [H3O+] with a term we call the pH of a solution. At 25 °C pH > 7 is basic pH < 7 is acidic pH = 7 is neutral pH = - log [H3O+] 47 48 12 pH scale pOH and acidity At 25 °C pOH = - log [OH-] Or [OH-] = 10-pOH pOH < 7 is basic pOH > 7 is acidic pOH = 7 is neutral 49 50 Problem Kw = 1.0 x 10-14 = [H3O+] [OH-] pKw = - log (1.0 x 10-14) = 14.00 (2 sigfigs! The 14 is not significant!) pKw = - log ([H3O+] [OH-]) = (- log [H3O+]) + (- log [OH-]) = pH + pOH Calculate the pH of each of the following solutions: a) A sample of seawater that has an OH- concentration of 1.58 x 10-6 mol/L b) A sample of acid rain that has an H3O+ concentration of 6.0 x 10-5 mol/L so 14.00 = pH + pOH (at 25 °C)! 51 52 13 Problem Measuring pH Calculate [H3O+] and [OH-] in each of the following solutions: a) Human blood (pH 7.40) b) A cola beverage (pH 2.8) We often measure the pH of a solution with a chemical acid-base indicator. Indicators are B-L acids (symbolized HIn) where the acid form has a different colour than the conjugate base form (In-) HIn (aq) + H2O (l) ' H3O+ (aq) + In- (aq) colour A colour B 53 54 Measuring pH Indicators tend to change colour only in small pH ranges of about 2 units. 55 To make a universal indicator that covers the pH range from about 1 to 12, a mixture of several different indicators with different pH ranges is used. 56 14 Universal indicator Methyl violet Phenolphthalein Bromothymol blue Bromocresol green Universal indicator Methyl orange pH 1 2 3 4 5 6 7 8 9 10 11 57 58 15
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