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Chapter 16

Chapter 16. Acids and Bases. Overview. Acid – Base Concepts Arrhenius Br ø nsted – Lowry Lewis Acid and Base Strengths Relative Strengths of Acids and Bases Molecular Structure and Acid Strength Self – Ionization of Water and pH Self – Ionization of Water

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Chapter 16

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  1. Chapter 16 Acids and Bases

  2. Overview • Acid – Base Concepts • Arrhenius • Brønsted – Lowry • Lewis • Acid and Base Strengths • Relative Strengths of Acids and Bases • Molecular Structure and Acid Strength • Self – Ionization of Water and pH • Self – Ionization of Water • Solutions of a Strong Acid or Base • The pH of a Solution

  3. Acid – base concepts • Arhennius Concept: • acid: provides H+ ions (called a proton) in water. HA(aq)  H+(aq) + A(aq) • base: provides OH in water. • Brønsted–Lowry Concept: • acid: donates H+. • base: accepts H+. • No water or hydroxide required. • Base converted to its conjugate acid and develops properties that are those of an acid. E.g. Identify each as either an acid or base and determine its conjugate: and , F, H2CO3 and E.g. Identify the conjugate acids and bases in the reaction below: • Amphoteric substance: a substance that can act as either an acid or base. (E.g. H2O in examples)

  4. Acid – base concepts 2 • Lewis acid: electron pair acceptor • Lewis base: electron pair donor.. • Lewis definition most general, then Brønsted-Lowery and finally Arhennius: • E.g. • E.g.2 determine the Lewis acid and base in BF3 and NH3 in the reaction: • E.g.3 determine the Lewis acid and base in the following reaction: Co3+(aq) + 6F(aq)  [CoF6]3.

  5. Acid/Base Strength • Extent of acid/base reaction variable and depends upon the relative strengths of the acids and conjugate bases.. • Strong acids and bases react with other reactant to produce all product. Stronger acids and bases react to form weaker conjugate bases and acids.

  6. Relative Strengths of Acids and Bases and Extent of Reaction • The table of relative strengths of acids and conjugate bases can be used to predict if a reaction will produce product. E.g. Which will produce product? HNO3 + CNor HCN +

  7. Factors Affecting Acid Strength • Binary acids: • Bond strength is directly related to the acid strength (bond size). • HI and HBr have larger bonds lengths and are more acidic than HF and HCl, even though fluorine is most electronegative. • For bonds of similar size the acid strength is related to electronegativity difference. E.g. bonds across a row more acidic towards right side of periodic table. • Oxyacids are acidic substances that contain oxygen and some other nonmetal, e.g. HNO3, HOCl, etc.. Anything that affects the polarity of the O–H bond will affect the strength of the acid. • An increase in the electronegativity of an atom bound to oxygen increases in polarity of the bond and makes it more acidic. • More oxygen = more polar. E.g. determine relative acidity of HOI, HOBr and HOCl. E.g. HClO4 is strongest acid in its oxyacid series.

  8. Autoionization of Water • Water can act as both an acid and a base equilibrium is: H2O + H2O  H3O+ + OH. Kw = [OH][H3O+] = 1.00 x 1014M2. • Since [OH] = [H3O+][H3O+] = 1 x 107 M (called a neutral solution) • Acidic [H3O+] > 1.00x107M • Neutral [H3O+] = 1.00x107M • Basic [H3O+] < 1.00x107M • All acids/bases dissolved in water must obey equation for the ionization of water. • They either add H3O+ or OH to water. • Most of the acids in this chapter will be stronger than water and add significantly to the hydronium ion concentration. E.g. the hydronium ion concentration of an acidic solution was 1.00x105 M. What was the [OH]? • E.g. what is the hydronium ion concentration if the hydroxide concentration was 2.50x103 M?

  9. Strong acids and Bases • A strong acid is completely dissociated in water. This leads to [H3O+] and [OH]. Eg: Calculate the [H3O+],[OH] and [Cl] for a 0.048 M HCl solution. Assume the contribution from water is negligible. E.g.2 Calculate the conc. of all ionic species as well as the pH if CNaOH = 0.080 M. E.g. 3 what is pH and [OH] of 0.125 M Ba(OH)2. • [H3O+] of water is small compared to added [H3O+] from the acid and ignored in the calculation. • A more rigorous treatment is: • The last term is very small except when the concentration of the strong acid is very small. • When the concentration of added acid is small compared with pure water, include the contribution from water. E.g.4 estimate the [H3O+] of 108 M HCl.

  10. The pH of a Solution • pH = log[H3O+] and [H3O+] = 10pH • Acidic pH < 7.00 • Neutral pH = 7.00 • Basic pH > 7.00 E.g. determine the pH of a solution in which [H3O+] = 5.40x106 M E.g.2 determine the pH of a solution in which the [OH] = 3.33x103 M E.g.3 determine the pOH of a solution in which the [OH] = 3.33x103 M E.g.4 Determine the [H3O+] if the pH of the solution is 7.35. • The term pX is defined in exactly the same way as pH. Eg.5 What is the pCa if [Ca2+] = 6.44x10-4

  11. Methods of Measuring pH • pH paper is used that has compounds in it which are change to different colors for different pH ranges. • An colored indicator can be placed in the solution and its color correlated with pH. HIn(aq) + H2O(l)  H3O+(aq) + In(aq). E.g. phenolphthalein is colorless in acid form but pink in basic form. • The pH at which they change color depends on their equilibrium constant. • More accurate and precise measurements are made with a pH meter. A combination of voltmeter and electrodes.

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