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Chapter 3 Molecular Shape and Structure 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 CHAPTER 10 ACIDS AND BASES THE NATURE OF ACIDS AND BASES 10.1 BrØnsted-Lowry Acids and Bases In aqueous or nonaqueous solution, Acid : Proton donor Base : Proton acceptor +− HCl(aq) +⎯H23O(l) ⎯→+H O (aq) Cl (aq) Fig. 10.1 Dissolution of HCl in water. Proton-transfer reaction ⎯⎯→ +− HCN(aq) ++H23O(l)←⎯⎯ H O (aq) CN (aq) 2+ − Ca (aq) ++HCO32(aq) H O(l) ⎯⎯→ + ←⎯⎯ H33O (aq) + CaCO (s) A strong acid is fully deprotonated. A weak acid is only partially deprotonated in solution. Fig. 10.2 Stalactites(종유석,鍾乳石) and Stalagmites(석순,石筍) 2– ▷ Oxide, O , is a strong base. ▷ Ammonia, NH3, is a weak base. 2−− + − O(aq)+⎯H2O(l) ⎯→2 OH(aq) NH32(aq) ++H O(l) R NH4(aq) OH (aq) Fig. 10.3 An oxide ion in water Fig. 10.4 Equilibrium structure of aqueous ammonia solution 1 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 ▶ Conjugate acids and bases + Acid ⎯⎯donates⎯ H⎯→ Conjutgate base ⎯⎯→ + − HCN(aq) ++HO2 (l)←⎯⎯ H3O (aq) CN(aq) acid 1 base 2 acid 2 base 1 + Base ⎯accept⎯⎯⎯s H → Conjutgate acid + − NH(3 aq)++HO24()laR NH ( q) OH(aq) base 1 acid 2 acid 1 base 2 Ex. 10.1 Writing the formulas of conjugate acids and bases − 2− (a) The conjugate base of HCO3 is CO3 2− − Æ The conjugate acid of CO3 is HCO3 . (b) The conjugate base of OH− is O2− . Æ The conjugate acid of O2− is OH− . ☺ Brønsted-Lowry definition applies to nonaqueous solvents −+ CH33COOH(l) ++NH (l) R CH3CO2(am) NH4(am) ☺ Brønsted-Lowry definition applies also to the gas phase. HCl(gg) +⎯NH34( ) ⎯→ NH Cl(s) Fig. 10.5 White powder (NH4Cl) formed from NH3(g) and HCl(g) 10.2 Lewis Acids and Bases Most general definition: Lewis acid : Electron pair acceptor Lewis base : Electron pair donor 2− O is a Lewis base: NH3 is a Lewis base 2 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 Fig. 10.6 Actions of acids (left) and bases (right) in Arrhenius, BrØnsted, and Lewis definitions. Lewis acid (CO2) + Lewis base (H2O) Æ Brønsted acid (H2CO3) 10.3 Acidic, Basic, and Amphoteric Oxides ▶ Acidic oxide (CO2): Molecular compound (nonmetal), Reacts with water to form a Brønsted acid 2 NaOH(aq) + CO22(g) ⎯⎯→+Na CO3(aq) H2O(l) NaOH(aq) + CO23(g) ⎯⎯→ NaHCO (aq) ▶ Basic oxide (CaO, MgO): Ionic salt (metal) Reacts with water to form a Brønsted base Reacts with acid to form a salt and water CaO(sl) + H22O( ) ⎯⎯→Ca(OH) (aq) MgO(s) + 2 HCl(l) ⎯⎯→+MgCl22(aq) H O(l) ▶ Amphoteric oxide (Al2O3): Reacts with both acid and base Al23O (sl) + 6 HCl( ) ⎯⎯→+2 AlCl3(aq) 3 H2O(l) 2 NaOH(aq) + Al O (s) + 3 H O(l) → 2 Na[Al(OH) ](aq) 23 2 4 Fig. 10.7 Amphoteric oxides of metalloid elements. 3 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 Fig. 10.8 The d-block elements forming amphoteric oxides with intermediate oxidation states. 10.4 Proton Exchange Between Water Molecules ▶ Amphiprotic (양쪽성양성자성): acting both as a proton donor and as a proton acceptor ZZX +− o −1 2 H23O(la)YZZ HO()q+ OH(aq), ∆=H r +56 kJ ⋅mol (A) Æ Autoprotolysis (자체양성자이전반응) or Autoionization aa+− HO3 OH K = 2 a ()HO2 Kaw = +a− : Autoprotolysis constant or water product constant HO3 OH o o1− aJ ≈ [J]/ c, where c =⋅1 mol L +− −77− −14o Kw ==[H3O ][OH ] ()1.0 ×10 ×(1.0 ×10 ) =1.0×10 at 25 C + Fig. 10.9 Autoprotolysis of water Fig. 10.10 The product of the concentrations of hydronium (H3O ) and hydroxide (OH–) ions in water is a constant. 4 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 Ex. 10.2 Calculating the concentrations of ions in a solution of a metal hydroxide. o + − A solution of 0.0030 M Ba(OH)2(aq) at 25 C. [H3O ] = ? [OH ] = ? Decide whether the Ba, alkaline earth metal compound is fully Æ Ba(OH)2 dissociates almost completely dissociated in solution. in water to produce OH– ions – 2+− Find the mole ratio of [OH ] Ba(OH)2 (sa) ⎯⎯→+Ba ( q) 2OH (aq) – to solute concentration. 1 mol Ba(OH)2 ≈ 2 mol OH Calculate [OH–] from the [OH−−] = 2 ×⋅0.0030 mol L 1 solute concentration. =⋅0.0060 mol L−1 Rearrange −14 + Kw 1.0 ×10 +− [H3O ] ==− Kw = [H3O ][OH ] [OH ] 0.0060 −−12 1 + =×1.7 10 mol ⋅L to find [H3O ]. 10.5 The pH Scale + pH =−log a + , pH =−log[H3O ] HO3 pH < 7 (acidic) pH =−log[1.0 ×10−7 ] =7.00 (pure water at 25oC) pH > 7 (basic) Ex. 10.3 Calculating a pH from a concentration. + −−81 (a) What is the pH of human blood in which [H3O ] = 4.0×⋅10 mol L ? + −8 pH =−log[H3O ] pH =−log[4.0 ×10 ] =7.40 (b) 0.020 M HCl(aq) + pH =−log[H3O ] pH =−log0.020 =1.70 (c) 0.040 M KOH(aq) [OH−−] ==[KOH] 0.040 mol ⋅L 1 K 1.0 ×10−14 [H O+−] ==w =2.5×10 13 3 [OH − ] 0.040 + −13 pH =−log[H3O ] pH =−log(2.5×10 ) =12.60 5 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 Fig. 10.11 A pH-meter is a voltmeter measuring the pH electronically. (a) Orange juice (b) Lemon juice Fig. 10.12 Typical pH values of common aqueous solutions. 10.6 The pOH of Solutions pXl=− ogX pOH =−loga OH− pOH =−log[OH − ] −14 pKKww=−log =−log(1.0×10 ) =14.00 + − log Kw = log[H3O ][OH ] +− −−log[H3O ] log[OH ] =−log Kw Fig. 10.13 The pH and pOH scales. Always, pH + pOH = 14 pH +=pOH pKw =14 6 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 WEAK ACIDS AND BASES Fig. 10.14 Magnesium metal dissolved in (a) HCl and (b) HAc with the same concentrations. The rate of hydrogen evolution depends on the concentration of hydronium ions. 10.7 Acidity and Basicity Constants ▶ Molecular composition of a solution of a weak acid (or base) - acid (or base) molecules or ions + – - small concentrations of H3O (or OH ) ions and the conjugate base of the acid – + - very small concentrations of OH (or H3O ) ions maintaining autoprotolysis equilibrium ⎯⎯→ + − CH32COOH(aq) ++H O(l)←⎯⎯ H3O (aq) CH3CO2(aq) aa+− K = HO3 CH32CO aa CH32COOH H O + − [H33O ][CH CO2] −5 o Ka ==1.8×10 at 25 C [CH3COOH] ⎯⎯→ +− HA()aq ++H23O(l)←⎯⎯ HO()aq A(aq) +− [H3O ][A ] Ka = , plKKaa=− og [HA] Fig. 10.15 A solution of a weak acid. ⎯⎯→ + − NH32(aq) ++H O(l)←⎯⎯ NH4(aq) OH (aq) + − aa+− NH OH [NH ][OH ] K = 4 Æ K =4 =1.8×10−5 at 25oC , p4K = .75 aa b [NH ] b NH32H O 3 7 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 + − ⎯⎯→ +− [BH ][OH ] B(aq) ++H 2O(l) ←⎯⎯ BH (aq) OH (aq) , Kb = , pKKbb= −log [B] 10.8 The Conjugate Seasaw ▶ Relative strength of an acid (or a base) and its conjugate base (or acid): Æ the stronger the acid (or base), the weaker its conjugate base (or acid) + ⎯⎯→ + NH42()aq ++HO(l)←⎯⎯ H3O()aq NH3(aq), + [H33O ][NH ] Ka = + [NH4 ] [H O+ ][ NH ] [ NH + ][OH− ] KK = 3 3 × 4 = [H O+ ][OH− ] ab + [ NH ] 3 [ NH4 ] 3 KKab×=Kw (11a) log KKab+=log log Kw pKKab+=p pKw (11b) + pKKaw(NH43) =−p pKb(NH ) =14.00 −4.75 =9.25 Fig. 10.17 Conjugate acid-base pairs. 8 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 Ex. 10.5 Deciding which of two species is the stronger acid or base. (a) HF / HIO3 (b) HNO2 / HCN Ka(HIO3) > Ka(HF) Ka(HNO2) > Ka(HCN) pKa(HIO3)=0.77 < pKa(HF)=3.45 pKa(HNO2)=3.37 < pKa(HCN)=9.31 10.9 Molecular Structure and Acid Strength Binary acid, HA The greater the electronegativity of A, the stronger the acid HA. The acid strengths increase (while bond strengths decrease) down the group. The more polar the H-A bond, the stronger the acid across a row. 9 2009년도 제2학기 화 학 2 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 10 10.10 The Strengths of Oxoacids and Carboxylic Acids Proton of an –OH group in an oxoacid is acidic due to high polarity of the O–H bond. ▶ Phosphorus acid, H3PO3 Æ (HO)2PHO Æ Donates H’s from –OH groups Æ No donation of H from –PH group Electronegativity difference Æ between O and P ▶ Oxoacids with the different central halogen atoms but with the same number of O atoms Æ The greater the electronegativity of the halogen, the stronger the oxoacid. ▶ Oxoacids with the same central atom with the different number of O atoms Æ The greater the number of O atoms attached to the central atom, the stronger the oxoacid. Æ The greater the oxidation number of the central atom, the stronger the oxoacid.
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