Ch17 Acid Base A

Information about Ch17 Acid Base A

Published on January 3, 2008

Author: Brainy007

Source: authorstream.com

Content

The Chemistry of Acids and Bases:  The Chemistry of Acids and Bases Chapter 17 Acid and Bases:  Acid and Bases Acid and Bases:  Acid and Bases Acid and Bases:  Acid and Bases Strong and Weak Acids/Bases:  Strong and Weak Acids/Bases Generally divide acids and bases into STRONG or WEAK ones. STRONG ACID: HNO3(aq) + H2O(liq) ---> H3O+(aq) + NO3-(aq) HNO3 is about 100% dissociated in water. Slide6:  HNO3, HCl, H2SO4 and HClO4 are among the only known strong acids. Strong and Weak Acids/Bases Slide7:  Weak acids are much less than 100% ionized in water. One of the best known is acetic acid = CH3CO2H Strong and Weak Acids/Bases Slide8:  Strong Base: 100% dissociated in water. NaOH(aq) ---> Na+(aq) + OH-(aq) Strong and Weak Acids/Bases Other common strong bases include KOH and Ca(OH)2. CaO (lime) + H2O --> Ca(OH)2 (slaked lime) Slide9:  Weak base: less than 100% ionized in water One of the best known weak bases is ammonia NH3(aq) + H2O(liq) e NH4+(aq) + OH-(aq) Strong and Weak Acids/Bases ACID-BASE THEORIES:  ACID-BASE THEORIES The most general theory for common aqueous acids and bases is the BRØNSTED - LOWRY theory ACIDS DONATE H+ IONS BASES ACCEPT H+ IONS ACID-BASE THEORIES:  The Brønsted definition means NH3 is a BASE in water — and water is itself an ACID ACID-BASE THEORIES ACID-BASE THEORIES:  ACID-BASE THEORIES NH3 is a BASE in water — and water is itself an ACID NH3 / NH4+ is a conjugate pair — related by the gain or loss of H+ Every acid has a conjugate base - and vice-versa. Conjugate Pairs:  Conjugate Pairs More About Water:  More About Water H2O can function as both an ACID and a BASE. In pure water there can be AUTOIONIZATION Equilibrium constant for autoion = Kw Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC More About Water:  More About Water Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC In a neutral solution [H3O+] = [OH-] so Kw = [H3O+]2 = [OH-]2 and so [H3O+] = [OH-] = 1.00 x 10-7 M Autoionization Calculating [H3O+] & [OH-]:  Calculating [H3O+] & [OH-] You add 0.0010 mol of NaOH to 1.0 L of pure water. Calculate [H3O+] and [OH-]. Solution 2 H2O(liq) e H3O+(aq) + OH-(aq) Le Chatelier predicts equilibrium shifts to the ____________. [H3O+] < 10-7 at equilibrium. Set up a ICE table. Calculating [H3O+] & [OH-]:  Calculating [H3O+] & [OH-] You add 0.0010 mol of NaOH to 1.0 L of pure water. Calculate [H3O+] and [OH-]. Solution 2 H2O(liq) e H3O+(aq) + OH-(aq) initial 0 0.0010 change +x +x equilib x 0.0010 + x Kw = (x) (0.0010 + x) Because x << 0.0010 M, assume [OH-] = 0.0010 M [H3O+] = Kw / 0.0010 = 1.0 x 10-11 M Calculating [H3O+] & [OH-]:  Calculating [H3O+] & [OH-] You add 0.0010 mol of NaOH to 1.0 L of pure water. Calculate [H3O+] and [OH-]. Solution 2 H2O(liq) e H3O+(aq) + OH-(aq) [H3O+] = Kw / 0.0010 = 1.0 x 10-11 M This solution is _________ because [H3O+] < [OH-] [H3O+], [OH-] and pH:  [H3O+], [OH-] and pH A common way to express acidity and basicity is with pH pH = log (1/ [H3O+]) = - log [H3O+] In a neutral solution, [H3O+] = [OH-] = 1.00 x 10-7 at 25 oC pH = -log (1.00 x 10-7) = - (-7) = 7 [H3O+], [OH-] and pH:  [H3O+], [OH-] and pH What is the pH of the 0.0010 M NaOH solution? [H3O+] = 1.0 x 10-11 M pH = - log (1.0 x 10-11) = 11.00 General conclusion — Basic solution pH > 7 Neutral pH = 7 Acidic solution pH < 7 Slide21:  Figure 17.1 [H3O+], [OH-] and pH:  [H3O+], [OH-] and pH If the pH of Coke is 3.12, it is ____________. Because pH = - log [H3O+] then log [H3O+] = - pH Take antilog and get [H3O+] = 10-pH [H3O+] = 10-3.12 = 7.6 x 10-4 M pH of Common Substances:  pH of Common Substances Other pX Scales:  Other pX Scales In general pX = -log X and so pOH = - log [OH-] Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC Take the log of both sides -log (10-14) = - log [H3O+] + (-log [OH-]) pKw = 14 = pH + pOH Equilibria Involving Weak Acids and Bases:  Equilibria Involving Weak Acids and Bases Aspirin is a good example of a weak acid, Ka = 3.2 x 10-4 Equilibria Involving Weak Acids and Bases:  Equilibria Involving Weak Acids and Bases Acid Conjugate Base acetic, CH3CO2H CH3CO2-, acetate ammonium, NH4+ NH3, ammonia bicarbonate, HCO3- CO32-, carbonate A weak acid (or base) is one that ionizes to a VERY small extent (< 5%). Equilibria Involving Weak Acids and Bases:  Equilibria Involving Weak Acids and Bases Consider acetic acid, CH3CO2H (HOAc) HOAc + H2O e H3O+ + OAc- Acid Conj. base (K is designated Ka for ACID) Because [H3O+] and [OAc-] are SMALL, Ka << 1. Equilibrium Constants for Weak Acids:  Equilibrium Constants for Weak Acids Weak acid has Ka < 1 Leads to small [H3O+] and a pH of 2 - 7 Equilibrium Constants for Weak Bases:  Equilibrium Constants for Weak Bases Weak base has Kb < 1 Leads to small [OH-] and a pH of 12 - 7 Ionization Constants for Acids/Bases :  Ionization Constants for Acids/Bases Acids Conjugate Bases Increase strength Increase strength Relation of Ka, Kb, [H3O+] and pH:  Relation of Ka, Kb, [H3O+] and pH K and Acid-Base Reactions :  K and Acid-Base Reactions Reactions always go from the stronger A-B pair (larger K) to the weaker A-B pair (smaller K). K and Acid-Base Reactions:  K and Acid-Base Reactions A strong acid is 100% dissociated. Therefore, a STRONG ACID—a good H+ donor—must have a WEAK CONJUGATE BASE—a poor H+ acceptor. HNO3(aq) + H2O(liq) e H3O+(aq) + NO3-(aq) STRONG A base acid weak B Every A-B reaction has two acids and two bases. Equilibrium always lies toward the weaker pair. Here K is very large. K and Acid-Base Reactions:  K and Acid-Base Reactions We know from experiment that HNO3 is a strong acid. 1. It is a stronger acid than H3O+ 2. H2O is a stronger base than NO3- 3. K for this reaction is large K and Acid-Base Reactions:  K and Acid-Base Reactions Acetic acid is only 0.42% ionized when [HOAc] = 1.0 M. It is a WEAK ACID HOAc + H2O e H3O+ + OAc- WEAK A base acid STRONG B Because [H3O+] is small, this must mean 1. H3O+ is a stronger acid than HOAc 2. OAc- is a stronger base than H2O 3. K for this reaction is small Types of Acid/Base Reactions:  Types of Acid/Base Reactions Strong acid + Strong base H+ + Cl- + Na+ + OH- e H2O + Na+ + Cl- Net ionic equation H+(aq) + OH-(aq) e H2O(liq) K = 1/Kw = 1 x 1014 Mixing equal molar quantities of a strong acid and strong base produces a neutral solution. Types of Acid/Base Reactions:  Types of Acid/Base Reactions Weak acid + Strong base CH3CO2H + OH- e H2O + CH3CO2- This is the reverse of the reaction of CH3CO2- (conjugate base) with H2O. OH- stronger base than CH3CO2- K = 1/Kb = 5.6 x 104 Mixing equal molar quantities of a weak acid and strong base produces the acid’s conjugate base. The solution is basic. Types of Acid/Base Reactions:  Types of Acid/Base Reactions Strong acid + Weak base H3O+ + NH3 e H2O + NH4+ This is the reverse of the reaction of NH4+ (conjugate acid of NH3) with H2O. H3O+ stronger acid than NH4+ K = 1/Ka = 5.6 x 104 Mixing equal molar quantities of a strong acid and weak base produces the bases’s conjugate acid. The solution is acid. Types of Acid/Base Reactions:  Types of Acid/Base Reactions Weak acid + Weak base Product cation = conjugate acid of weak base. Product anion = conjugate base of weak acid. pH of solution depends on relative strengths of cation and anion. Types of Acid/Base Reactions: Summary:  Types of Acid/Base Reactions: Summary Calculations with Equilibrium Constants:  Calculations with Equilibrium Constants pH of an acetic acid solution. What are your observations? 0.0001 M 0.003 M 0.06 M 2.0 M a pH meter, Screen 17.9 Equilibria Involving A Weak Acid:  Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 1. Define equilibrium concs. in ICE table. [HOAc] [H3O+] [OAc-] initial change equilib Equilibria Involving A Weak Acid:  Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 1. Define equilibrium concs. in ICE table. [HOAc] [H3O+] [OAc-] initial 1.00 0 0 change -x +x +x equilib 1.00-x x x Note that we neglect [H3O+] from H2O. Equilibria Involving A Weak Acid:  Equilibria Involving A Weak Acid Step 2. Write Ka expression You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. This is a quadratic. Solve using quadratic formula or method of approximations (see Appendix A). Equilibria Involving A Weak Acid:  Equilibria Involving A Weak Acid Step 3. Solve Ka expression You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. First assume x is very small because Ka is so small. Now we can more easily solve this approximate expression. Equilibria Involving A Weak Acid:  Equilibria Involving A Weak Acid Step 3. Solve Ka approximate expression You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. x = [H3O+] = [OAc-] = [Ka • 1.00]1/2 x = [H3O+] = [OAc-] = 4.2 x 10-3 M pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37 Equilibria Involving A Weak Acid:  Equilibria Involving A Weak Acid Consider the approximate expression For many weak acids [H3O+] = [conj. base] = [Ka • Co]1/2 where C0 = initial conc. of acid Useful Rule of Thumb: If 100•Ka < Co, then [H3O+] = [Ka•Co]1/2 Equilibria Involving A Weak Acid:  Equilibria Involving A Weak Acid Calculate the pH of a 0.0010 M solution of formic acid, HCO2H. HCO2H + H2O e HCO2- + H3O+ Ka = 1.8 x 10-4 Approximate solution [H3O+] = [Ka • Co]1/2 = 4.2 x 10-4 M, pH = 3.37 Exact Solution [H3O+] = [HCO2-] = 3.4 x 10-4 M [HCO2H] = 0.0010 - 3.4 x 10-4 = 0.0007 M pH = 3.47 Weak Bases:  Weak Bases Equilibria Involving A Weak Base:  Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O e NH4+ + OH- Kb = 1.8 x 10-5 Step 1. Define equilibrium concs. in ICE table [NH3] [NH4+] [OH-] initial change equilib Equilibria Involving A Weak Base:  Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O e NH4+ + OH- Kb = 1.8 x 10-5 Step 1. Define equilibrium concs. in ICE table [NH3] [NH4+] [OH-] initial 0.010 0 0 change -x +x +x equilib 0.010 - x x x Equilibria Involving A Weak Base:  Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O e NH4+ + OH- Kb = 1.8 x 10-5 Step 2. Solve the equilibrium expression Assume x is small (100•Kb < Co), so x = [OH-] = [NH4+] = 4.2 x 10-4 M and [NH3] = 0.010 - 4.2 x 10-4 ≈ 0.010 M The approximation is valid ! Equilibria Involving A Weak Base:  Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O e NH4+ + OH- Kb = 1.8 x 10-5 Step 3. Calculate pH [OH-] = 4.2 x 10-4 M so pOH = - log [OH-] = 3.37 Because pH + pOH = 14, pH = 10.63 Slide54:  MX + H2O ----> acidic or basic solution? Consider NH4Cl NH4Cl(aq) ----> NH4+(aq) + Cl-(aq) (a) Reaction of Cl- with H2O Cl- + H2O ----> HCl + OH- base acid acid base Cl- ion is a VERY weak base because its conjugate acid is strong. Therefore, Cl- ----> neutral solution Acid-Base Properties of Salts Slide55:  NH4Cl(aq) ----> NH4+(aq) + Cl-(aq) (b) Reaction of NH4+ with H2O NH4+ + H2O ----> NH3 + H3O+ acid base base acid NH4+ ion is a moderate acid because its conjugate base is weak. Therefore, NH4+ ----> acidic solution See TABLE 17.4 for a summary of acid-base properties of ions. Acid-Base Properties of Salts Acid-Base Properties of Salts :  Acid-Base Properties of Salts Slide57:  Calculate the pH of a 0.10 M solution of Na2CO3. Na+ + H2O ---> neutral CO32- + H2O e HCO3- + OH- base acid acid base Kb = 2.1 x 10-4 Step 1. Set up concentration table [CO32-] [HCO3-] [OH-] initial change equilib Acid-Base Properties of Salts Slide58:  Calculate the pH of a 0.10 M solution of Na2CO3. Na+ + H2O ---> neutral CO32- + H2O e HCO3- + OH- base acid acid base Kb = 2.1 x 10-4 Step 1. Set up ICE table [CO32-] [HCO3-] [OH-] initial 0.10 0 0 change -x +x +x equilib 0.10 - x x x Acid-Base Properties of Salts Slide59:  Calculate the pH of a 0.10 M solution of Na2CO3. Na+ + H2O ---> neutral CO32- + H2O e HCO3- + OH- base acid acid base Kb = 2.1 x 10-4 Acid-Base Properties of Salts Assume 0.10 - x ≈ 0.10, because 100•Kb < Co x = [HCO3-] = [OH-] = 0.0046 M Step 2. Solve the equilibrium expression Slide60:  Calculate the pH of a 0.10 M solution of Na2CO3. Na+ + H2O ---> neutral CO32- + H2O e HCO3- + OH- base acid acid base Kb = 2.1 x 10-4 Acid-Base Properties of Salts Step 3. Calculate the pH [OH-] = 0.0046 M pOH = - log [OH-] = 2.34 pH + pOH = 14, so pH = 11.66, and the solution is ________.

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