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Published on January 5, 2008

Author: Heather

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Chapter Fifteen:  Chapter Fifteen Acids, Bases, And Acid-Base Equilibria The Arrhenius Theory (section 2.8):  The Arrhenius Theory (section 2.8) acid – a substance that produces H+ ion when dissolved in water base - a substance that produces OH- ion when dissolved in water strong acid – ionizes essentially completely into H+ and an anion strong base – dissociates nearly completely into OH- and a cation weak acid and base – ionize reversibly Limitation of Arrhenius Theory The Brønsted-Lowry Theory:  The Brønsted-Lowry Theory Acid – proton donor Base – proton acceptor Conjugate acid and base, HA/A-, differ by one proton. The conjugate acid of a base is the base plus the attached proton and the conjugate base of an acid is the acid minus the proton. A substance that can act either as an acid or a base is amphiprotic. For weak acids and bases, equations can be written to describe equilibrium conditions. Ionization of Ammonia:  Ionization of Ammonia Strengths Of Conjugate Acid-Base Pairs:  Strengths Of Conjugate Acid-Base Pairs The stronger an acid, the weaker is its conjugate base. The stronger a base, the weaker is its conjugate acid. An acid-base reaction is favored in the direction from the stronger member to the weaker member of each conjugate acid-base pair. Ka and Kb values are used to compare the strengths of weak acids and bases. Water has a leveling effect; when the strong acids are dissolved in water, they all completely ionize to the hydronium ion. Relative Acid-Base Pair Strength:  Relative Acid-Base Pair Strength Slide7:  Leveling effect– between the dotted lines also weak acids and bases Strong acids – Weak conjugate bases Strong bases – Weak conjugate acids Acid And Base Ionization Constants:  Acid And Base Ionization Constants weak acid: CH3COOH + H2O º H3O+ + CH3COO- [H3O+][CH3COO-] Acid ionization constant: Ka = [CH3COOH] weak base: NH3 + H2O º NH4+ + OH- [NH4+][OH-] Base ionization constant: Kb = [NH3] Acid and base ionization constants are the measure of the strengths of acids and bases. Relative Strengths Of Binary Acids:  Relative Strengths Of Binary Acids H –X The greater the tendency for the transfer of a proton from HX to H2O, the more the forward reaction is favored and the stronger the acid. in a periodic group: Bond-dissociation energy is inversely proportional to acid strength. The weaker the bond, the stronger the acid. Anion radius is directly proportional to acid strength. The larger the resultant anion’s radius, the stronger is the acid. The strengths of binary acids increase from top to bottom in a group of the periodic table. Relative Strengths Of Binary Acids:  Relative Strengths Of Binary Acids H –X in a periodic group: Bond dissociation energy: the weaker the bond, the stronger the acid. Bond dissociation energy 569 > 431 > 368 > 297 (kJ/mol) HF HCl HBr HI Acid strength Ka 6.6x10-4 < ~106 < ~108 < ~109 Anion radius: the larger the anion’s radius, the stronger the acid. Anion radius (ppm) 136 < 181 < 195 < 216 (kJ/mol) HF HCl HBr HI Acid strength Ka 6.6x10-4 < ~106 < ~108 < ~109 The strength of binary acids increase from top to bottom in a group of the periodic table. Relative Strengths Of Binary Acids:  Relative Strengths Of Binary Acids H –X in a period: The larger the electronegativity difference between H and X, the more easily the proton is removed and the stronger is the acid.  EN 0.4 < 0.9 < 1.4 < 1.9 Acid strength CH4 NH3 H2O HF The strengths of binary acids increase from left to right across a period of the periodic table. Representative Trends In Strengths of Binary Acids:  Representative Trends In Strengths of Binary Acids Strengths Of Oxoacids:  Strengths Of Oxoacids H – O - E Two factors: - electronegativity of the central atom (E) - number of terminal oxygen atoms As the electronegativity of the central atom (E) increases and as the number of terminal oxygen atoms increases, the acid strength also increases. Strengths Of Oxoacids:  Strengths Of Oxoacids As the electronegativity of the central atom (E) increases the acid strength increases. Electronegativity 2.5 < 2.8 < 3.0 HOI HOBr HOCl Acid strength Ka 2.3x10-11 < 2.5x10-9 < 2.9x10-8 As the number of terminal oxygen atoms increases, the acid strength also increases. # of terminal 0 1 2 3 O atoms O O H-O-Cl H-O-Cl-O H-O-Cl-O H-O-Cl-O O Acid strength 2.9x10-8 < 1.1x10-2 < ~1000 < ~108 Strengths Of Carboxylic Acids:  Strengths Of Carboxylic Acids O R – C – O - H Carboxylic acids all have the -COOH group in common; therefore, differences in acid strength must come from differences in the R group attached to the carboxyl group. In general, the more that electronegative atoms are attached in the R group, the stronger the acid. Strengths Of Carboxylic Acids:  Strengths Of Carboxylic Acids In general, the more that electronegative atoms are attached in the R group, the stronger the acid. I-CH2CH2COOH Cl-CH2CH2COOH CH3-CHClCOOH CH3CCl2COOH Ka 8.3x10-5 < 1.0x10-4 < 1.4x10-3 < 8.7x10-3 Strengths Of Amines As Bases:  Strengths Of Amines As Bases Aromatic amines are much weaker bases than aliphatic amines. This is due in part to the fact that the p electrons in the benzene ring of an aromatic molecule are delocalized and can involve the N’s lone-pair electrons in the resonance hybrid. As a result, the lone-pair electrons are much less likely to accept a proton. Electron-withdrawing groups on the ring further diminish the basicity of aromatic amines relative to aniline. Strengths Of Amines As Bases:  Strengths Of Amines As Bases BrNH2 NH3 C6H5NH2 Kb 2.5x10-8 1.8x10-5 7.4x10-10 Slide19:  Amine bases: N H H aromatic amine base R= CH3, CH2CH3 aliphatic amine bases ammonia Self-Ionization Of Water:  Self-Ionization Of Water Even the purest of water conducts electricity. This is due to the fact that water self-ionizes, that is, it creates a small amount of H3O+ and OH-. H2O + H2O º H3O+ + OH- Kw = [H3O+][OH-] Kw - ion product of water Kw = 1.0 x 10-14 at 25 oC This equilibrium constant is very important because it applies to all aqueous solutions - acids, bases, salts, and non-electrolytes - not just to pure water. Slide21:  O H H O H H O Self ionization reaction of water: + H + + - pH and pOH:  pH and pOH pH = - log[H3O+] [H3O+] = 10-pH pH = - log[OH-] [OH-] = 10-pOH pKw = pH + pOH = 14.00 neutral solution: [H3O+] = [OH-] = 10 –7 M pH = 7.0 acidic solution: [H3O+] > 10-7 M pH < 7.0 basic solution: [H3O+] < 10-7 M pH > 7.0 The pH Scale:  The pH Scale An Example:  An Example The pH of milk of magnesia, a suspension of solid magnesium hydroxide in its saturated aqueous solution, is measured to be 10.52. What is the molarity of Mg(OH)2 in its saturated aqueous solution? Equilibrium In Solutions Of Weak Acids And Weak Bases:  Equilibrium In Solutions Of Weak Acids And Weak Bases weak acid: HA + H2O º H3O+ + A- [H3O+][A-] Ka = [HA] weak base: B + H2O º HB+ + OH- [HB+][OH-] Kb = [B] You need to be able to write acid and base ionization equations!!! pKa and pKb:  pKa and pKb pKa = -logKa pKb = -logKb larger Ka => smaller pKa => stronger acid larger Kb => smaller pKb => stronger base Some Acid-Base Equilibrium Calculations:  Some Acid-Base Equilibrium Calculations These calculations are similar to the equilibrium calculations performed in Chapter 14. An equation is written for the reversible reaction, data are organized under this equation, the changes that occur in establishing equilibrium are assessed, and finally calculations of equilibrium concentrations are done. When Macid/Ka > 100 or Mbase/Kb > 100, the calculations can be simplified. An Example :  An Example 1.Determine the concentrations of H3O+, CH3COOH and CH3COO-, and the pH of 1.00 M CH3COOH solution. Ka = 1.8 x 10-5. 2. What is the pH of a solution that is 0.200 M in methylamine, CH3NH2? Kb = 4.2 x 10-4. Are Salts Neutral, Acidic or Basic?:  Are Salts Neutral, Acidic or Basic? Salts are ionic compounds formed in the reaction between an acid and a base. 1. NaCl Na+ is from NaOH , a strong base Cl- is from HCl, a strong acid H2O NaCl (s) > Na+ (aq) + Cl- (aq) Na+ and Cl- ions do not react with water. The solution is neutral. Are Salts Neutral, Acidic or Basic?:  Are Salts Neutral, Acidic or Basic? 2. KCN K+ is from KOH , a strong base CN- is from HCN, a weak acid H2O KCN (s) > K+ (aq) + CN- (aq) K+ ions do not react with water, but CN- ions do. CN- + H2O º HCN + OH- hydrolysis The OH- ions are produced, so the solution is basic. Are Salts Neutral, Acidic or Basic?:  Are Salts Neutral, Acidic or Basic? 3. NH4Cl NH4+ is from NH3 , a weak base Cl- is from HCl, a strong acid H2O NH4Cl (s) > NH4+ (aq) + Cl- (aq) Cl- ions do not react with water, but NH4+ ions do. NH4+ + H2O º H3O+ + NH3 hydrolysis The H3O+ ions are produced, so the solution is acdic. Hydrolysis:  Hydrolysis The hydrolysis of an ion is the reaction of an ion with water to produce the conjugate acid and hydroxide ion or the conjugate base and hydrogen ion. You need to be able to write equation for hydrolysis reaction! Ions As Acids And Bases:  Ions As Acids And Bases Certain ions can cause an aqueous solution to become acidic or basic due to hydrolysis. Salts of strong acids and strong bases form neutral solutions. Salts of weak acids and strong bases form basic solutions. Salts of strong acids and weak bases form acidic solutions. Salts of weak acids and weak bases form solutions that are acidic in some cases, neutral or basic in others. Strong Acids And Strong Bases:  Strong Acids And Strong Bases Table 4.1, p.141 Strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4 Strong bases: Group IA and IIA hydroxides - Memorize!! An Example:  An Example Indicate whether the solutions (a) Na2S and (b) KClO4 are acidic, basic or neutral. The pH Of A Salt Solution:  The pH Of A Salt Solution What is the pH of 1.0 M NaCN solution? Hydrolysis of CN- ions: CN- + H2O º HCN + OH- CN- is a conjugate base of HCN. Ka of HCN can be found. What is Kb for CN-? Ka x Kb = Kw so, Kb = Kw/Ka Common Ion Effect Illustrated:  Common Ion Effect Illustrated 1.00 M CH3COOH 1.00 M CH3COOH – 1.00 M CH3COONa yellow: pH < 3.0 blue-violet: pH > 4.6 The Common Ion Effect:  The Common Ion Effect If one solution contains a weak acid and another contains the same acid and its conjugate base as a second solute, the two solutions have different pH values. The solution containing both the weak acid and its conjugate base has a pH much higher than the solution containing only the weak acid. The conjugate base is referred to as a common ion because it is found in both the weak acid and the anion. The common ion effect is the suppression of the ionization of a weak acid or a weak base by the presence of a common ion from a strong electrolyte. The Common Ion Effect - An Example:  The Common Ion Effect - An Example Calculate the pH of 1.00 M CH3COOH-1.00 M CH3COONa solution. Depicting Buffer Action:  Depicting Buffer Action Buffer Solutions:  Buffer Solutions A buffer solution is a solution that changes pH only slightly when small amounts of a strong acid or a strong base are added. A buffer contains a weak acid with its salt (conjugate base) or a weak base with its salt (conjugate acid) CH3COOH/CH3COONa NH3/NH4Cl How A Buffer Solution Works:  How A Buffer Solution Works The acid component of the buffer can neutralize small added amounts of OH-, and the basic component can neutralize small added amounts of H3O+. Pure water does not buffer at all. Henderson-Hasselbalch Equation For Buff Solutions:  Henderson-Hasselbalch Equation For Buff Solutions [conjugate base] pH = pKa + log [weak acid] If [weak acid] = [conjugate base], pH = pKa Requirements: The ratio of [conjugate base] to [weak acid] is between 0.10 and 10 [conjugate base]/Ka > 100, [weak acid]/Ka > 100 Calculations in Buffer Solutions:  Example 15.17: A buffer solution is 0.24 M NH3 and 0.20 M NH4Cl. What is the pH of this buffer? If 0.0050 mol NaOH is added to 0.500 L of this solution, what will be the pH? Example 15.18: What concentration of acetate ion in 0.500 M CH3COOH produces a buffer solution with pH = 5.00? Calculations in Buffer Solutions Buffer Capacity And Buffer Range:  Buffer Capacity And Buffer Range There is a limit to the capacity of a buffer solution to neutralize added acid or base, and this limit is reached before all of one of the buffer components has been consumed. In general, the more concentrated the buffer components in a solution, the more added acid or base the solution can neutralize. As a rule, a buffer is most effective if the concentrations of the buffer acid and its conjugate base are equal. Acid-Base Indicators:  Acid-Base Indicators An acid-base indicator is a weak acid having one color and the conjugate base of the acid having a different color. One of the “colors” may be colorless. HIn + H2O º H3O+ + In- color 1 color 2 Acid-base indicators are often used for applications in which a precise pH reading isn’t necessary. A common indicator used in introductory chemistry laboratories is litmus. Several Common Indicators:  Several Common Indicators Types Of Calculations In Acid-Base Equilibria:  Types Of Calculations In Acid-Base Equilibria pH, pOH – Kw = [H3O+][OH-]; pKw = pH + pOH Weak acid and weak base A salt aqueous solution – hydrolysis, KaKb = Kw Common ion effect Buffer solution – Henderson–Hasselbalch equation Neutralization Reactions:  Neutralization Reactions Neutralization is the reaction of an acid and a base. Titration is a common technique for conducting a neutralization. At the equivalence point in a titration, the acid and base have been brought together in exact stoichiometric proportions. The point in the titration at which the indicator changes color is called the end point. The indicator endpoint and the equivalence point for a neutralization reaction can be best matched by plotting a titration curve, a graph of pH versus volume of titrant. In a typical titration, 50 mL or less of titrant that is 1 M or less is used. Neutralization Reaction Strong Acid - Strong Base:  Example 15.20 Calculate the pH at the following points in the titration of 20.00 mL of 0.500 M HCl with 0.500 M NaOH. Before the addition of any NaOH (initial pH). After the addition of 10.00 mL of 0.500 M NaOH (half-neutralization point). After the addition of 20.00 mL of 0.500 M NaOH (equivalence point). After the addition of 20.20 mL of 0.500 M NaOH (beyond the equivalence point). Neutralization Reaction Strong Acid - Strong Base Titration Curve For Strong Acid - Strong Base:  Titration Curve For Strong Acid - Strong Base Features Of Titration Curve For Strong Acid - Strong Base:  Features Of Titration Curve For Strong Acid - Strong Base pH is low at the beginning. pH changes slowly until just before equivalence point. pH changes sharply around equivalence point. pH = 7.0 at equivalence point. Further beyond equivalence point, pH changes slowly. Any indicator whose color changes in pH range of 4 – 10 can be used in titration. Neutralization Reaction Weak Acid - Strong Base:  Example 15.21 Calculate the pH at the following points in the titration of 20.00 mL of 0.500 M CH3COOH with 0.500 M NaOH. Before the addition of any NaOH (initial pH). After the addition of 8.00 mL of 0.500 M NaOH (buffer region). After the addition of 10.00 mL of 0.500 M NaOH (half-neutralization point). After the addition of 20.00 mL of 0.500 M NaOH (equivalence point). After the addition of 21.00 mL of 0.500 M NaOH (beyond the equivalence point). Neutralization Reaction Weak Acid - Strong Base Titration Curve For Weak Acid - Strong Base:  Titration Curve For Weak Acid - Strong Base Features Of Titration Curve For Weak Acid - Strong Base:  Features Of Titration Curve For Weak Acid - Strong Base The initial pH is higher because weak acid is partially ionized. At the half-neutralization point, pH = pKa. pH is greater than 7 at equivalence point because the anion of the weak acid hydrolyzes. The steep portion of titration curve around equivalence point has a smaller pH range. The choice of indicators for the titration is more limited. Types Of Calculations In Acid-Base Equilibria:  Types Of Calculations In Acid-Base Equilibria pH, pOH – Kw = [H3O+][OH-]; pKw = pH + pOH Weak acid and weak base A salt aqueous solution – hydrolysis, KaKb = Kw Common ion effect Buffer solution – Henderson–Hasselbalch equation Neutralization, titration curve Lewis Acids And Bases:  Lewis Acids And Bases There are reactions in non-aqueous solvents, in the gaseous state, and even in the solid state that can be considered acid-base reactions in which Brønsted-Lowry theory is not adequate to explain. A Lewis acid is a species that is an electron-pair acceptor and a Lewis base is a species that is an electron-pair donor. In organic chemistry, Lewis acids are often called electrophiles and Lewis bases are often called nucleophiles. Summary:  Summary In the Brønsted-Lowry theory an acid is a proton donor and a base is a proton acceptor. If an acid is strong, its conjugate base is weak; and if a base is strong, its conjugate acid is weak. Water is amphiprotic: it can be either an acid or a base. It undergoes limited self-ionization producing H3O+ and OH-. pH = -log[H3O+] pOH = -log[OH-] pKw = -logKw The pH in both pure water and in neutral solutions is 7. Acidic solutions have a pH less than 7 and basic solutions have a pH greater than 7. Summary (continued):  Summary (continued) In aqueous solutions at 25 oC, pH + pOH = 14.00. Hydrolysis reactions cause certain salt solutions to be either acidic or basic. A strong electrolyte that produces an ion common to the ionization equilibrium of a weak acid or a weak base suppresses the ionization of the weak electrolyte. Acid-base indicators are weak acids for which the acid and its conjugate base have different colors. In Lewis acid-base theory, a Lewis acid accepts an electron pair and a Lewis base donates an electron pair.

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