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Acids and Bases pH and Titrations

Acids and Bases pH and Titrations. 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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Acids and Bases pH and Titrations

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  1. Acids and BasespH and Titrations

  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. II. Theories of acid and bases A. Arrhenius theory 1. Arrhenius Acid as known as a traditional acid a. a chemical compound that contains hydrogens and ionizes in aqueous solution to form hydrogen ions 2. Arrhenius base a. any chemical that produces hydroxide ions

  4. B. Bronsted -Lowry theory 1923 (worked independently) 1. Bronsted Acid a. an ion or molecule that is a proton donor 2. Bronsted base a. anything which will accept a proton

  5. 3. Conjugate acid - base pairs HF + HOH ===> H3O+ + F- Acid1 Base2 Acid2 Base1 a. Strong acids have weak conjugate bases b. Strong bases and weak conjugate acids 4. Autoionization - self-ionization HOH <===> H+ + OH -

  6. Conjugate Acid-Base Pairs Fig. 17-1, p. 507

  7. Conjugate Acid-Base Pairs p. 507

  8. Brønsted-Lowry Theory of Acids & Bases Conjugate Acid-Base Pairs General Equation p. 507

  9. Brønsted-Lowry Theory of Acids & Bases p. 504

  10. 5. Amphiprotic ability to act as an acid or base C. Lewis theory 1. Lewis acid a. anything which can accept a pair of electrons 2. Lewis Base a. anything which will donate a pair of electrons

  11. Lewis Theory of Acids & Bases p. 506

  12. Objectives • Properties of acids and bases • The pH scale • Distinguish between strong and weak acids and list the clinical uses of these acids • Distinguish between strong and weak bases and list the clinical uses of these acids • Understand neutralisation and the clinical applications of neutralisation

  13. On page 355 Do questions 4,6,7,10,14, 17,18,22,23,25,26,28

  14. Properties of Acids Aqueous solutions of acids have a sour taste. Acids change the color of acid-base indicators. Some acids react with active metals to release hydrogen gas. Acids react with bases to produce salts and water. Some acids conduct electric current.

  15. PROPERTIES OF ACIDS & BASES Acids Produce hydrogen ions (H+) in H2O Taste sour Act as electrolytes in solution Neutralise solutions containing hydroxide ions (OH -) React with several metals releasing H2(g)  corrosion React with carbonates releasing CO2(g) Destroy body tissue

  16. How many foods can you think of that are sour? Chances are, almost all of the foods that you associate with being sour, owe their sour taste to an acid: Lemons – citric acid Grapefruit – citric acid Apples – malic acid Sour milk – lactic acid Vinegar – acetic acid Grapes – tartaric acid

  17. Strength of acids Astrong acid is one that ionizes completely in aqueous solutions. Acids that are weak electrolytes are weak acids. The strength of an acid depends on the polarity of the bond between hydrogen and the element to which it is bonded and the ease with which that bond can be broken. Acid strength increases with increasing polarity and decreasing bond energy.

  18. Strong acids are: Strong electrolytes ~ 100% ionisation  good conductors Severe burns to body tissue *** Stomach lining protected against HCl by mucus

  19. Marieb, Fig 26.11 STRENGTHS OF ACIDS • Strong Acids (very few) Eg • HCl Hydrochloric Acid ~ Stomach acid • HNO3 Nitric Acid ~ May be used to cauterise warts ~Drugs, explosives, fertilisers, dyes • H2SO4 Sulphuric Acid ~  conc. to treat stomach hypoacidity ~ Fertilisers, dyes, glues

  20. Factors Affecting Acid Strength Binary acids: For bonds of similar size the acid strength is related to electronegativity difference. 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.

  21. Oxyacids are acidic substances that contain • oxygen and some other nonmetal, e.g. HNO3 • 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.

  22. Weak Acids (most acids in nature) CH3COOH Acetic Acid ~ Vaginal jellies, antimicrobial solution  ears, plastics, dyes, insecticides H2CO3 Carbonic Acid ~Bicarbonate buffer system, carbonated drinks H3PO4 Phosphoric Acid ~ Drugs, fertilisers, soaps, detergents, animal feed

  23. Produce or cause an increase in hydroxide ions (OH-) in H2O Bases Taste bitter Turn red litmus  blue Act as electrolytes in solution Neutralise solutions containing hydrogen ions (H +) Have a slippery, ‘soapy’ feel Destroy body tissue/ dissolve fatty (lipid) material

  24. How many bases can you think of? • Ammonia • Sodium hydroxide – lye – drain cleaner • Milk of magnesia – Mg(OH)2 – antacid • Aluminum hydroxide – antacid • Baking soda – sodium hydrogen carbonate

  25. 4. STRENGTHS OF BASES Strong Bases NaOH Sodium Hydroxide ~ Removes grease – drains, ovens Mg(OH)2 Magnesium hydroxide ~ Antacid ~ Laxative Al(OH)3 Aluminium hydroxide ~ Antacid ~ Absorbs toxins, gases, ~ Causes constipation

  26. Strength of Bases • The strength of a base depends on the extent to which the base dissociates. • Strong bases are strong electrolytes. • Strong bases: calcium hydroxide, barium hydroxide, sodium hydroxide, etc • Weak bases: ammonia, aniline • Table 15-4 page 461

  27. 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 +

  28. Strong bases are: • Strong electrolytes • ~ 100% dissociation in water  good conductors • Severe damage to skin & eyes (Group 1A elements)

  29. Weak Bases Eg NH3 Ammonia ~ Waste product of protein break down in body. CO3 2- In antacids HCO3 – In antacids, buffers HPO4 2- In buffers

  30. Weak bases are: Weak electrolytes Do not contain OH – but react with H2O  small numbers of OH – Reaction with Water : Weak bases NH3(g) + H2O NH4+(aq) + OH –(aq) HCO3 – (aq) + H2OH2CO3 (aq) + OH-(aq)

  31. Acids & Bases STRONG vs WEAK _ completely ionized _ partially ionized _ strong electrolyte _ weak electrolyte _ ionic/very polar bonds _ some covalent bonds Strong Acids:Strong Bases: HClO4LiOH H2SO4NaOH HIKOH HBrCa(OH)2 HClSr(OH)2 HNO3Ba(OH)2

  32. H2O NaOH(s) Na+(aq) + OH-(aq) • Dissociation in Water : Strong bases Metal hydroxides  ions H2O Mg(OH)2(s) Mg 2+ (aq) + OH- (aq) H2O Al(OH)3(s) Al 3+(aq) + OH-(aq)

  33. 5. ACID-BASE NEUTRALISATION • Neutralisation Reaction Acid + Base  Salt + Water HCl + NaOH  NaCl + H2O H+ + OH – H2O Neutralise each other Must be equal concentrations

  34. Antacids – clinical applications (Check for side effects!!) ~ Neutraliseexcess stomach acid ~ Raise stomach pH > 4 • Pepsin inactive ~ Assist with ulcer treatment ~  solubility in H2O but still produce high % of ions

  35. CaCO3 2HCl + CaCO3  CaCl2 + H2O + CO 2 (g) ~Also a Ca 2 + supplement Long term overuse  Ca 2 + levels •  risk kidney stones (renal calculi)

  36. Eg • Mg(OH) 2 Milk of Magnesia & in Mylanta 2HCl + Mg(OH) 2  MgCl2 + 2H2O • Al(OH) 3 In Mylanta 3HCl + Al(OH) 3  AlCl3 + 3H2O

  37. NaHCO3 Baking Soda Not recommended!! HCl + NaHCO3  NaCl + H2O + CO2 (g) ~ Elderly tend to OD  Stomach can ‘explode’

  38. Weak acids are: • Weak electrolytes • Small % ionisation  weak conductors • Dissociation in Water : Weak acids Polar covalent molecules  Mainly stay as molecules

  39. Dissociation in Water : Strong acids Polar covalent molecules  ions Eg. HCl(l) H+(aq) + Cl-(aq) HNO3(l) H+ (aq) + NO3- (aq) H2SO4(l) 2H+ (aq) + SO42-(aq) H2O H2O H2O

  40. H2O • Dissociation in water : Weak acids (cont) CH3COOH (l)H+(aq) + CH3COO-(aq) H2CO3 (l) H+ (aq) + HCO3-(aq) H2O H2O H3PO4 (l) H+(aq) + H2PO4- (aq)

  41. 2. THE pH SCALE Ion Product of Water Pure H2O at 25°C Some molecules ionise H2O H+ + OH- [H+ ] = 1 x 10-7 M = [OH- ]

  42. On the pH scale, values below 7 are acidic, a value of 7 is neutral, and values above 7 are basic.

  43. The pH scale The pH scale ranges from 1x 100 to 1 x 10-14 mol/L or from 1 to 14. pH = - log [H+] 1 2 3 4 5 678 9 10 11 12 13 14 acid neutral base

  44. Stomach juice: pH = 1.0 – 3.0 Human blood: pH = 7.3 – 7.5 Lemon juice: pH = 2.2 – 2.4 Seawater: pH = 7.8 – 8.3 Vinegar: pH = 2.4 – 3.4 Ammonia: pH = 10.5 – 11.5 Carbonated drinks: pH = 2.0 – 4.0 0.1M Na2CO3: pH = 11.7 Orange juice: pH = 3.0 – 4.0 1.0M NaOH: pH = 14.0

  45. pH describes [H+ ] & [OH- ] • Indicates if a fluid is : 0 Acidic [H+ ] = 100[OH- ] =10-14 7 Neutral [H+ ] = 10-7[OH- ] =10-7 14 Basic [H+ ] = 10-14[OH- ] = 100

  46. Using the pH scale Exponential values for [H+ ] & [OH- ] inconvenient in a clinical workplace • Simplify  pH scale  acid-base concentration p  potential or Power H  Hydrogen

  47. Acidic solution [H+ ] > [OH- ] Neutral solution [H+ ] = [OH- ] Basic solution [H+ ] < [OH- ]

  48. Water Equilibrium Kw = [H+] [OH-] = 1.0 x 10-14 Equilibrium constant for water • Water or water solutions in which [H+] = [OH-] = 10-7 M are neutral solutions. • A solution in which [H+] > [OH-] is acidic • A solution in which [H+] < [OH-] is basic

  49. Autoionization of Water • Water can act as both an acid and a baseequilibrium 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)

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