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Properties of acids

Properties of acids. Electrolytes: conduct electricity React to form salts Change the color of an indicator Have a sour taste. Properties of Bases. Bitter taste “slippery feel” Electrolyte React with an acid to form a salt. Acid names:. ide=  hydro______ic acid

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Properties of acids

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  1. Properties of acids • Electrolytes: conduct electricity • React to form salts • Change the color of an indicator • Have a sour taste

  2. Properties of Bases • Bitter taste • “slippery feel” • Electrolyte • React with an acid to form a salt

  3. Acid names: • ide= hydro______ic acid • ite == ____________ous acid • ate == _____________ic acid

  4. Arrhenius Acids and Bases • An acid contains a Hydrogen ion that easily disassociates • A base has a hydroxide ion that easily disassociates

  5. Bronsted-Lowry • An acid is a hydrogen ion donor • A base is a hydrogen ion acceptor. • NH3(aq)+ H2O(aq)  NH4+(aq) + OH-(aq) • Base Acid Conj. Acid Conj. Base

  6. conjugate acids and bases • These are the reverse reaction “reactants” • HA(aq) + H2O(l) H3O+(aq) + A-(aq) • Acid Base conjugate conjugate acid base • HCl + NaOH  HOH + Cl-

  7. Lewis acid and bases • An acid is an electron pair acceptor • A base is an electron pair donator base acid • NH3 +BF3 H3N BF3

  8. Acid Strength A strong acid is one for which the equilibrium lies far to the products side Ka > 1. A weak acid is one for which the equilibrium lies far to the reactants side Ka < 1. Table 14.1 on page 659.

  9. Monoprotic: HCl • Diprotic: H2SO4 • Triprotic H3PO4

  10. Oxyacids • Acids where the acidic hydrogen is attached to an oxygen. • Pictures on pages 658-659. • Structural formulas given in table 14.8 on page 694.

  11. Organic acids • Acids with a carbon atom backbone. • Commonly contain a carboxyl group. • Acetic acid • Benzoic acid

  12. Amphoteric Substance • A substance that can act as both an acid and a base. • Water • Ammonia

  13. We can tell an acid from a base by using an indicator

  14. Autoionization of water • H2O + H2O  H3O+ + OH- • Kw = [H3O+] [OH-] = 1.0 x 10-14 • [H3O+] [OH-] = [H+] [OH-] • [H+] = [OH-] =1.0 x 10-7 M

  15. pH = -log [H+] • [H+] = antilog (-pH) • pOH = -log[OH-] • pK = -log K • Significant figures in logarithms- the # of decimal places in the log = # of sig figs in the original number.

  16. [H+] [OH-] = 1.0 x 10-14 • pH + pOH = 14 • For strong acids the [H+] = the molarity of the acid.

  17. System for solving weak acid equilibrium problems • List the major species in solution. • Find any species that can produce H+ and write a balanced equation for the reaction producing H+. • Use the values for K for the reactions you have written to decide which reaction will dominate.

  18. Write the equilibrium expression for the dominant reaction. • List the initial concentrations of the species in the dominant reaction. • Define the change needed to obtain equilibrium. Define x.

  19. Write equilibrium concentrations in terms of x. • Put equilibrium concentrations into equilibrium expression. • Solve for x the “easy way” • Use the 5% rule to see if approximation is valid. • Calculate [H+] and pH.

  20. Calculate the pH of a 0.100M solution of hypochlorous acid.

  21. Calculate the pH of a solution that contains 1.00M HCN and 1.00M HNO2. Also calculate the concentration of cyanide ions.

  22. % dissociation • = amount dissociated x100 • initial concentration • For a given weak acid, the % dissociated increases as the acid becomes more dilute.

  23. In a 0.100M solution, lactic acid (HC3H5O3)is 3.7% dissociated. Calculate the value of Ka for this acid.

  24. Bases • B(aq) + H2O(l)  HB+(aq) + OH-(aq) • Base acid conj conj • acid base • Kb = [BH+][OH-] • [B]

  25. Bases • Strong bases – hydroxides of group 1A metals and calcium, barium and strontium. • Weak bases are commonly ammonia and substituted ammonia compounds. • Table 14.3 on page 678.

  26. Calculate the pH for a 15.0M solution of NH3. Kb = 1.8 x10-5

  27. Salts that produce neutral solutions • Salts that consist of the cations of strong bases and the anions of strong acids have no effect on pH. • KCl, NaNO3 , Ba(HSO4)2

  28. Salts that produce basic solutions • Salts that consist of cations of strong bases and the anion is the conjugate base of a weak acid. • NaC2H3O2, KNO2, Sr(CN)2

  29. Salts that produce acidic solutions • Salts that consist of a cation that is the conjugate acid of a weak base and the anion of a strong acid. • NH4Cl • A salt that contains a highly charged metal ion. AlCl3 see sample exercise 14.20.

  30. For any weak acid-conjugate base or weak base-conjugate acid • Ka x Kb = Kw

  31. Calculate the pH of: • 0.10M NaCl • 0.10M NaF • 0.10M NH4Cl

  32. If the anion is the conjugate base of a weak acid and the cation is the conjugate acid of a weak base the Kb must be compared to the Ka. Which ever is greater will dominate.

  33. Predict if the following will be acidic or basic • NH4C2H3O2 • NH4CN • NH4NO2

  34. Effect of structure on acid-base properties Acidic properties depend on two factors: H-X • The strength of the bond . As the strength increases the acidity decreases • The polarity of the bond. As polarity increases acidity increases

  35. Strength of oxyacids • Within a series of oxyacids as the number of oxygens increases the strength of the acid also increases. • Table 14.8 page 694

  36. Acid-base properties of oxides • Depends on the electronegativity of the element bonded to oxygen. O-X • Non-metal oxides in water will form acids. O-X is stronger than H-O in polar water. • Metal oxides in water form bases. O-H stronger than O-X in polar water.

  37. Common Ion Effect • The shift in equilibrium position because of the addition of an ion already involved in the equilibrium. • Equilibrium shifts away from the added component.

  38. Buffered Solutions • Resists change in pH when either hydrogen or hydroxide ions are added. • Consist of a solution that contains both a weak acid and its salt or a weak base and its salt. Important Characteristics are on page 726.

  39. Buffer Capacity • Represents the amount of hydrogen or hydroxide ions the buffer can absorb without a significant change in pH. • Best buffers contain a weak acid with a pKa as close as possible to the desired pH.

  40. Indicators • A substance that changes color depending on the pH of the solution it is in. • When choosing an indicator we want the indicator end point and the titration equivalence point to be as close as possible.

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