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Acids and Bases

Acids and Bases

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Acids and Bases

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  1. Acids and Bases Original: L. Scheffler Modified: Swiftney

  2. 8.1 Theories of acids and bases 2 hours Original: L. Scheffler Modified: Swiftney

  3. 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-). Categorical definition – easy to sort substances into acids and bases Problem – many bases do not actually contain hydroxides Original: L. Scheffler Modified: Swiftney

  4. 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- Ex. HCl + H2O  H3O+ + Cl- Acid Base Conj Acid Conj Base • Operational definition - The classification depends on how the substance behaves in a chemical reaction. Original: L. Scheffler Modified: Swiftney

  5. 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 base/acid pair • * :A- - conjugate base of acid HA • * :B & HB+ - conjugate acid/base pair • * HB+ - conjugate acid of base :B Original: L. Scheffler Modified: Swiftney

  6. Examples of Bronsted-Lowry Acid Base Systems • * Note: Water can act as acid or base (Amphoteric) • 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- Original: L. Scheffler Modified: Swiftney

  7. G.N. Lewis Definition • Lewis • Acid - an electron pair acceptor • Base - an electron pair donor • *Note: form dative bonding Original: L. Scheffler Modified: Swiftney

  8. Acid Base Definition Summary • Acid Base • Arrhenius: H+ OH- (based on H2O) • B-L: H+ donor H+ acceptor (e.g. NH3) • Lewis: Electron Pair Electron Pair • Acceptor Donor Original: L. Scheffler Modified: Swiftney

  9. 8.2 Properties of acids and bases 1 hour Original: L. Scheffler Modified: Swiftney

  10. 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 aceticacid. Citrus fruits contain citricacid. Original: L. Scheffler Modified: Swiftney

  11. Acids • React with certain metals to produce hydrogen gas. • 2HCl + 2Na  2NaCl + H2 • React with carbonates and bicarbonates to produce carbondioxide gas • MgCO3 + 2HCl  MgCl2 + H2CO3  H2O + CO2 Bases • Have a bittertaste • Feel slippery. • Many soaps contain bases. Original: L. Scheffler Modified: Swiftney

  12. Properties of Acids • Produce H+ (as 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 “Blue to Red A-CID” (BRA) Original: L. Scheffler Modified: Swiftney

  13. 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 “BasicBlue” Original: L. Scheffler Modified: Swiftney

  14. 8.3 Strong and weak acids and bases 2 hours Original: L. Scheffler Modified: Swiftney

  15. Acid Base Dissociation • Acid-base reactions are equilibrium processes. • The ratio of the concentrations of the reactants and products is constant for a given temperature at equilibrium • It is known as the Acid or Base Dissociation Constant. • The stronger the acid or base, the larger the value of the dissociation constant. Original: L. Scheffler Modified: Swiftney

  16. Acid Strength • Strong Acid - Transfers all of its protons to water; - Completely ionized (dissociated); - Strong electrolyte; - The conjugate base is weaker and has a negligible tendency to be protonated. • Weak Acid - Transfers only a fraction of its protons to water; • - Partially ionized (dissociated); - Weak electrolyte; - The conjugate base is stronger, readily accepting protons from water • 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 Original: L. Scheffler Modified: Swiftney

  17. Acid Dissociation Constants Dissociation constants for some weak acids Original: L. Scheffler Modified: Swiftney

  18. Base Strength • Strong Base - all molecules accept a proton; - Completely ionized (dissociated); - strong electrolyte; - conjugate acid is very weak, negligible tendency to donate protons. • Weak Base - fraction of molecules accept proton; - Partially ionized (dissociated); - weak electrolyte; - the conjugate acid is stronger. It more readily donates protons. • 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. Original: L. Scheffler Modified: Swiftney

  19. Common Strong Acids/Bases Strong Acids Hydrochloric Acid Nitric Acid Sulfuric Acid Perchloric Acid Strong Bases Sodium Hydroxide Potassium Hydroxide *Barium Hydroxide *Calcium Hydroxide *While strong bases they are not very soluble Original: L. Scheffler Modified: Swiftney

  20. 8.4 The pH scale 1 hour Original: L. Scheffler Modified: Swiftney

  21. Acid Base Equilibrium Strong acid H3O+(aq) + A-(aq) HA(aq) 0.1 M 0.1M 0.1 M e.g. HCl, HNO3 assume 100% dissociation Weak Acid HA(aq) H3O+((aq) + A-(aq) e.g. CH3COOH <<0.1 M <<0.1 M 0.1 M pH = - log [H3O+] ***The pH scale is used to describe the concentration of acid present in a solution pH is used to make an “ugly” number into something simple (between 0-14). Original: L. Scheffler Modified: Swiftney

  22. The pH Scale * pH = - log [H3O+] pH [H3O+ ] [OH- ] pOH Original: L. Scheffler Modified: Swiftney

  23. pH and acidity pH = - log [H3O+] orpH = - log [H+] 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. Original: L. Scheffler Modified: Swiftney

  24. pH of Common Substances Original: L. Scheffler Modified: Swiftney

  25. 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 [H+] (or [OH-]) as a means of avoiding using large numbers and powers. * pH = - log [H3O+] = log(1 / [H3O+]) * pOH = - log [OH-] = log(1 / OH-]) • In pure water, [H3O+] = 1 x 10-7 mol / L (at 20oC)  pH = - log(1 x 10-7)= - (0 - 7) = 7 • pH range of solutions: 0 - 14 pH < 7 (Acidic) [H3O+]; pH > 7 (Basic) [H3O+] Original: L. Scheffler Modified: Swiftney

  26. 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 Original: L. Scheffler Modified: Swiftney

  27. pH Practice 1) pH of a solution containing 25 g HCl dissolved in 1.5 L of H2O? 2) pH of a solution containing 1.32 g of HNO3 dissolved in 750 mL? 3) pH of a solution containing 1.2 moles of Nitric Acid and 1.7 moles of Hydrochloric acid dissolved in 1000 mL? 4) If a solution has a [H+] of 4.5 x 10-7 M, is this acidic or basic? Explain Original: L. Scheffler Modified: Swiftney

  28. pH and acidity * Kw = [H3O+] [OH-] = 1.0 x10-14 In pure water * [H3O+] = [OH-] = 1.0 x10-7 * pH + pOH = 14 Original: L. Scheffler Modified: Swiftney

  29. Neutralization An acid will neutralize a base, giving a salt and water as products (THIS IS STOICHIOMETRY) Examples Acid Base Salt water HCl + NaOH  NaCl + H2O H2SO4 + 2 NaOH  Na2SO4 + 2 H2O H3PO4 + 3 KOH  K3PO4 + 3 H2O 2 HCl + Ca(OH) 2 CaCl2 + 2 H2O A salt is an ionic compound that is formed from thepositive ion (cation)of thebaseand thenegative ion (anion) of the acid Original: L. Scheffler Modified: Swiftney

  30. Neutralization Calculations If the concentration of acid or base is expressed in Molarity or mol dm-3 then: --The volume in dm3 multiplied by the concentration yields moles (mol) . -- If the volume is expressed in cm3 the same product yields millimoles (mmol) mol dm-3 x dm3 = mole * mol dm-3 x cm3 = (0.001) x mole = mmol Original: L. Scheffler Modified: Swiftney

  31. Neutralization Problems • The volume of solution in dm3 multiplied by concentration in moles dm-3 will yield moles. • If an acid and a base combine in a 1 to 1 ratio, the moles of acid will equal the moles of base. *** This is a DIFFERENT way of doing Acid/Base Stoichiometry • Therefore the volume of the acid multiplied by the concentration of the acid is equal to the volume of the base multiplied by the concentration of the base. VacidC acid = V base C base If any three of the variables are known, it is possible to determine the fourth. Original: L. Scheffler Modified: Swiftney

  32. Neutralization Problems Example 1:Hydrochloric acid reacts with potassium hydroxide according to the following reaction: HCl + KOH  KCl + H2O If 15.00 cm3 of 0.500 M HCl exactly neutralizes 24.00 cm3 of KOH solution, what is the concentration of the KOH solution? Solution: Vacid Cacid = Vbase Cbase (15.00 cm3 )(0.500 M) = (24.00 cm3 ) Cbase Cbase = (15.00 cm3 )(0.500 M) (24.00 cm3 ) Cbase = 0.313 M Original: L. Scheffler Modified: Swiftney

  33. Neutralization Problems Whenever an acid and a base do not combine in a 1 to 1 ratio, a mole factor must be added to the neutralization equation n Vacid C acid = V base C base The mole factor (n) is the number of times the moles the acid side of the above equation must be multiplied so as to equal the base side. (or vice versa) Example H2SO4 + 2 NaOH  Na2SO4 + 2 H2O The mole factor is 2 and goes on the acid side of the equation. The number of moles of H2SO4 is one half that of NaOH. Therefore the moles of H2SO4 are multiplied by 2 to equal the moles of NaOH. Original: L. Scheffler Modified: Swiftney

  34. Neutralization Problems Example 2:Sulfuric acid reacts with sodium hydroxide according to the following reaction: H2SO4 + 2 NaOH  Na2SO4 + 2 H2O If 20.00 cm3 of 0.400 M H2SO4 exactly neutralizes 32.00 cm3 of NaOH solution, what is the concentration of the NaOH solution? Solution: In this case the mole factor is 2 and it goes on the acid side, since the mole ratio of acid to base is 1 to 2. Therefore 2 Vacid Cacid = Vbase Cbase 2 (20.00 cm3 )(0.400 M) = (32.00 cm3 ) Cbase Cbase = (2) (20.00 cm3 )(0.400 M) (32.00 cm3 ) Cbase = 0.500 M Original: L. Scheffler Modified: Swiftney

  35. Neutralization Problems Example 3:Phosphoric acid reacts with potassium hydroxide according to the following reaction: H3PO4 + 3 KOH  K3PO4 + 3 H2O If 30.00 cm3 of 0.300 M KOH exactly neutralizes 15.00 cm3 of H3PO4 solution, what is the concentration of the H3PO4 solution? Solution: In this case the mole factor is 3 and it goes on the acid side, since the mole ratio of acid to base is 1 to 2. Therefore 3 VacidCacid = VbaseCbase 3 (15.00 cm3 )(Cacid) = (30.00 cm3 ) (0.300 M) Cacid= (30.00 cm3 )(0.300 M) (3) (15.00 cm3 ) Cacid= 0.200 M Original: L. Scheffler Modified: Swiftney

  36. Neutralization Problems Example 4:Hydrochloric acid reacts with calcium hydroxide according to the following reaction: 2 HCl + Ca(OH)2 CaCl2 + 2 H2O If 25.00 cm3 of 0.400 M HCl exactly neutralizes 20.00 cm3 of Ca(OH)2 solution, what is the concentration of the Ca(OH)2 solution? Solution: In this case the mole factor is 2 and it goes on the base side, since the mole ratio of acid to base is 2 to 1. Therefore Vacid Cacid = 2 Vbase Cbase (25.00 cm3) (0.400) = (2) (20.00 cm3) (Cbase) Cbase = (25.00 cm3 ) (0.400 M) (2) (20.00 cm3 ) Cbase = 0.250 M Original: L. Scheffler Modified: Swiftney

  37. Weak Acid Equilibria A weak acid is only partially ionized. Both the ion form and the unionized form exist at equilibrium HA + H2O  H3O+ + A- The acid equilibrium constant is Ka = [H3O+ ] [A-] [HA] Ka values are relatively small for most weak acids. The greatest part of the weak acid is in the unionized form Original: L. Scheffler Modified: Swiftney

  38. Water as an Equilibrium System • Water has the ability to act as either a Bronsted- Lowry acid or base. • Autoionization – spontaneous formation of low concentrations of [H+] and OH-] ions by proton transfer from one molecule to another. • Equilibrium Constant for Water Original: L. Scheffler Modified: Swiftney

  39. Amphoteric Solutions • A chemical compound able to react with both an acid or a base is amphoteric.    • Water is amphoteric. The two acid-base couples of water are H3O+/H2O and H2O/OH-It behaves sometimes like an acid, for example • And sometimes like a base : • Hydrogen carbonate ion HCO3- is also amphoteric, it belongs to the two acid-base couples H2CO3/HCO3- and HCO3-/CO32- Original: L. Scheffler Modified: Swiftney