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Acids and Bases

Acids and Bases. Chapter 15. Acids and Bases. The concepts acids and bases were loosely defined as substances that change some properties of water. One of the criteria that was often used was taste. Substances were classified salty-tasting, sour-tasting, sweet-tasting, bitter-tasting

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Acids and Bases

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  1. Acids and Bases Chapter 15

  2. Acids and Bases • The concepts acids and bases were loosely defined as substances that change some properties of water. • One of the criteria that was often used was taste. • Substances were classified • salty-tasting, sour-tasting, sweet-tasting, bitter-tasting • Sour-tasting substances would give rise to the word 'acid', which is derived from the Greek word oxein, which mutated into the Latin verb acere, which means 'to make sour' • Vinegar is a solution of acetic acid. Citrus fruits contain citric acid.

  3. Acids • React with certain metals to produce hydrogen gas. • React with carbonates and bicarbonates to produce carbon dioxide gas Bases • Have a bitter taste • Feel slippery. • Many soaps contain bases.

  4. Properties of Acids • Produce H+ (or H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule) • Taste sour • Corrode metals • Good Electrolytes • React with bases to form a salt and water • pH is less than 7 • Turns blue litmus paper to red

  5. Properties of Bases • Generally produce OH- ions in water • Taste bitter, chalky • Are electrolytes • Feel soapy, slippery • React with acids to form salts and water • pH greater than 7 • Turns red litmus paper to blue

  6. Arrhenius Definition Arrhenius Acid - Substances in water that increase the concentration of hydrogen ions (H+). Base - Substances in water that increase concentration of hydroxide ions (OH-).

  7. Bronsted-Lowry Definition Acid - neutral molecule, anion, or cation that donates a proton. Base - neutral molecule, anion, or cation that accepts a proton. HA + :B HB+ + :A- HCl + H2O  H3O+ + Cl- Acid Base Conj Acid Conj Base

  8. Conjugate Acid Base Pairs • Conjugate Base - The species remaining after an acid has transferred its proton. • Conjugate Acid - The species produced after base has accepted a proton. • HA & :A- - conjugate acid/base pair • :A- - conjugate base of acid HA • :B & HB+ - conjugate acid/base pair • HB+ - conjugate acid of base :B

  9. Examples of Bronsted-Lowry Acid Base Systems • Note: Water can act as acid or base • Acid Base Conjugate AcidConjugate Base • HCl + H2O  H3O+ + Cl- • H2PO4- + H2O  H3O+ +HPO42- • NH4+ + H2O  H3O+ + NH3 • Base Acid Conjugate AcidConjugate Base : NH3 + H2O  NH4++ OH- • PO43- + H2O  HPO42- + OH-

  10. G.N. Lewis Definition • Lewis • Acid - an electron pair acceptor • Base - an electron pair donor

  11. pH and acidity • HC2H3O2 + H2O  H3O+ + C2H3O2- • NH3 + HCl  NH4+ + Cl- • HSO4- + HPO4-2  H2SO4 + PO4-3 • HSO4- + HPO4-2  SO4-2 + H2PO4-1

  12. pH and acidity The pH values of several common substances are shown at the right. Many common foods are weak acids Some medicines and many household cleaners are bases.

  13. The pH Scale pH [H3O+ ] [OH- ] pOH

  14. pH and acidity • Acidity or Acid Strength depends on Hydronium Ion Concentration [H3O+] • The pH system is a logarithmic representation of the Hydrogen Ion concentration (or OH-) as a means of avoiding using large numbers and powers. pH = - log [H3O+] pOH = - log [OH-] In pure water [H3O+] = 1 x 10-7 M  pH = - log(1 x 10-7) = 7 • pH range of solutions: 0 - 14 pH < 7 (Acidic) [H3O+] > 1 x 10-7 M pH > 7 (Basic) [H3O+] < 1 x 10-7 M

  15. pH and acidity Important Equations pH = - log [H3O+] [H3O+] = 10-pH pOH = - log [OH-] [OH-] = 10-pOH pH + pOH = 14 Kw = [H3O+] [OH-] = 1.0 x10-14 In pure water: [H3O+] = [OH-]= 1.0 x10-7

  16. Calculating the pH pH = - log [H3O+] Example 1: If [H3O+] = 1 X 10-10pH = - log 1 X 10-10 pH = - (- 10) pH = 10 Example 2: If [H3O+] = 1.8 X 10-5pH = - log 1.8 X 10-5 pH = - (- 4.74) pH = 4.74

  17. pH and acidity If a substance has a pH of 3.5, determine: [H3O+], [OH-], pOH and whether it is acidic, basic, or neutral. If a substance has an [OH-] of 8.7 x 10-5, determine: pH, pOH, [H3O+] and whether it is acidic, basic, or neutral.

  18. Indicators

  19. Acid Strength • Strong Acid - Transfers all of its protons to water; - Completely ionized or dissociated; - Strong electrolyte; - The conjugate base is very weak • Weak Acid - Transfers only a fraction of its protons to water; • - Only partly ionizes or dissociates; - Weak electrolyte; - The conjugate base is stronger • As acid strength decreases, base strength increases. • The stronger the acid, the weaker its conjugate base • The weaker the acid, the stronger its conjugate base

  20. Base Strength • Strong Base - all molecules accept a proton; - completely ionizes or dissociates; - strong electrolyte; - conjugate acid is very weak • Weak Base - fraction of molecules accept proton; - partly ionizes or dissociates - weak electrolyte; - the conjugate acid is stronger. • As base strength decreases, acid strength increases. • The stronger the base, the weaker its conjugate acid. • The weaker the base the stronger its conjugate acid.

  21. Common Strong Acids Strong Acids Hydrochloric Acid, HCl Hydrobromic Acid, HBr Hydroioidic Acid, HI Nitric Acid, HNO3 Sulfuric Acid, H2SO4 Chloric Acid, HClO3 Perchloric Acid, HClO4

  22. Dissociation of Strong Acids • Strong acids completely dissociation (broke down) to the ions that make them up. • HCl  H+1 + Cl-1 • HNO3  H+1 + NO3-1 • If you know the concentration or molarity of a strong acid, you also know the amount of H+ ions and can find the pH (-log [H+]) • 0.1 M HCl = 0.1 M H+, pH = -log 0.1 = 1.0

  23. Common Strong Bases Strong Bases Any Group 1 Hydroxide and any Group 2 Hydroxide below Mg. Sodium Hydroxide, NaOH Potassium Hydroxide, KOH *Barium Hydroxide, Ba(OH)2 *Calcium Hydroxide, Ca(OH)2 *While strong bases they are not very soluble

  24. Dissociation of Strong Bases • Like strong acids, strong bases also completely dissociate to the ions that make them up. • KOH  K+1 + OH-1 • Ca(OH)2  Ca+2 + 2 OH-1 • Just like with strong acids, if you know the molarity of a strong base, you can determine the pOH and therefore the pH. • 0.01 M NaOH = 0.01 M OH-, • pOH = -log 0.01 = 2.0 and pH = 12.0

  25. pH of Strong Acids and Strong Bases • Determine the pH of 0.25 M KOH. • Determine the pH of 0.0050 M H2SO4. • Determine the pH of 0.0075 M Ca(OH)2.

  26. Dissociation of Weak Acids • Weak acids only dissociate to a very small degree, most break down to ions less than 5%. So an equilibrium is established where there is a small amount of product (H+ and conjugate base) but most of the acid does not break down. • These acids are weak electrolytes versus strong acids are strong electrolytes.

  27. Dissociation of Weak Acids • To determine the pH of a weak acid, an “ICE” chart must be used. • I = Initial concentrations • C = Change • E = Equilibrium concentrations • Unless otherwise noted, the initial concentrations of the products are zero.

  28. Ka = [H3O+] [A−] [HA] HA(aq) + H2O(l)  A−(aq) + H3O+(aq) Dissociation Constants • For a generic weak acid dissociation, the equilibrium expression would be • This equilibrium constant is called the acid-dissociation constant, Ka.

  29. Dissociation of Weak Acids • Equilibrium expressions are written as products/reactants, and liquids and solids are left out of the expressions. • Ka are used to tell the relative strength of the acid. • You will be given the Ka value for each weak acid.

  30. Dissociation of Weak Acids • Determine the pH of 0.025 M HC2H3O2. Ka = 1.8 X 10-5.

  31. Dissociation of Weak Acids • Determine the pH of 1.0 M hypoiodous acid, HIO. Ka of HIO = 2.3 x 10-11.

  32. Determination of Ka for a Weak Acid • If you are given the initial concentration of the weak acid, and the equilibrium concentration of the conjugate base or H+, you can determine the Ka of the acid. • Remember the initial concentrations of both products are zero. • If you know the equilibrium concentration of the conjugate base or H+, you also know the value of “x” in the “C” step of the “ICE” chart.

  33. Determination of Ka for a Weak Acid • Determine the Ka of a weak acid when the initial concentration of HA as 0.015 M and the equilibrium concentration of H+ = 0.0035 M. • Equation: HA + H2O  A-1 + H3O+1

  34. Titration • Titration – technique for determining the unknown concentration of an acid or base. • Titration involves delivery (from a buret) of a measured volume of a solution of known concentration (the titrant, usually a base) into a volume of solution of unknown concentration (the analyte, usually an acid).

  35. Titration • Equivalence or stoichiometric point – point in a titration where enough titrant has been added to react exactly with the analyte. • Equivalence pt: moles of acid = moles of base

  36. Titration

  37. Titration • The equivalence point is often marked by an indicator (commonly phenolphthalein), which is added at the beginning of the titration. It changes color at (or just after) the equivalence point. • The point where the indicator actually changes color is called the endpoint of the titration.

  38. Titration

  39. Titration Calculations • 2.00 mL of an unknown concentration of HCl was titrated with 0.25 M NaOH. If 30.23 mL of NaOH was needed to reach the stoichiometric point, what is the concentration of the HCl? • Equation: HCl + NaOH  NaCl + H2O

  40. Titration Calculations • 2 drops of phenolphthalein is added to 10.0 mL of an 0.75 M of HCl. If 16.82 mL of NaOH is needed to completely react with all of the HCl, what is the concentration of the NaOH?

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