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Acids & bases

Acids & bases. Acids and Bases reactions occur in everyday life and are essential for understanding our world. How does pH value affect our environment?. Why is it important to monitor and maintain the pH of the water in aquariums, soil and our blood? What exactly is pH? How is it measured?.

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Acids & bases

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  1. Acids & bases

  2. Acids and Bases reactions occur in everyday life and are essential for understanding our world. How does pH value affect our environment?

  3. Why is it important to monitor and maintain the pH of the water in aquariums, soil and our blood? What exactly is pH? How is it measured?

  4. Milk of magnesia is a medicine that usually relieves uncomfortable gastrointestinal symptoms within 30 minutes and constipation within six hours.Why is the milk of magnesia an antacid?

  5. Keywords • Acidity • Basicity (Monoprotic, diprotic, triprotic) • Bronsted-Lowry Theory - Proton donor/acceptor • Acid-base Conjugate pair • Amphiprotic • Lewis Theory - Lone pair electrons - Dative/Coordinate bond

  6. What is an acid? A solution that contains __________ ions (protons). OLD THEORY Weak acid like ethanoic acid does not have the power to neutralise strong acid like sodium hydroxide.

  7. What is a base/alkali? • A base is a substance like __________and ______________that reacts with acid to form salt and water only. • An alkali is a soluble base which in solution produces ________ ions. • Most bases are insoluble in water. 3 soluble bases are NaO/NaOH, KO/KOH, CaO/Ca(OH)2 Both acids and alkalis are _____________.

  8. What causes acidity? • It is the __________________that give an acid its acidic • properties when they dissolve in water and ____________ • into ions. • E.g. HCl gas is a covalent compound. • When dissolves in water, it forms HCl acid which dissolciate • to form ions.

  9. What is basicity (proticity)? • Basicity • refers to the no.of _____ atoms in one molecule of acid that • can be replaced by a _________. • refers to the no. of _______that can be replaced by one • molecule of that acid. • E.g. HCl (monobasic), • H2SO4(dibasic), • H3PO4(tribasic)

  10. Bronsted-Lowry theory An acid is defined as a molecule or ion that acts as a proton __________. A base is defned as a molecule or ion that acts as a proton __________.

  11. Types of acids • Acids that have single proton to donate – ___________ (monobasic). E.g. HCl(aq), HNO3(aq), HNO2(aq) • Acids that have 2 protons to donate – __________ E.g. H2SO4(aq), H2SO3(aq), H2CO3(aq) • H3PO4(aq) is ___________.

  12. Hydrogen chloride gas dissolved in water (solvent) HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) The equation can be split into (i) HCl(aq) Cl-(aq) + H+(aq) acid conjugate base (ii) H2O(l) + H+(aq) H3O+(aq) base conjugate acid Acidic behaviour is a transfer reaction in different solvents.

  13. Acid-base conjugate pair CH3COOH(l) + H2O(l) H3O+ (aq) + CH3OOO-(aq) conjugate conjugate acid base acid base donates H+ donates H+ NH3(g) + H2O(l) NH4+(aq) + OH-(aq) Water is sometimes described as _______________ because it can accept or donate a proton.

  14. Competition between acid/base and its conjugate (i) HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) acid base conjugate acid conjugate base (ii) CH3COOH(l) + H2O(l) H3O+ (aq) + CH3OOO-(aq) acid base conjugate acid conjugate base • Water is a much stronger base than chloride ion and has a stronger tendency to accept proton.The equilibrium shifts more to the right. • Ethanote ion is a much stronger base than water molecule. The equilbrium shifts to the left.

  15. Gas-phase acid-base reaction HCl(g) + NH3(g) NH4Cl(s) • The Bonsted-Lowry model can be extended to gas-phase acid-base reaction. • It involves the transfer of hydrogen ion from hydrogen chloride to ammonia.

  16. Strong acids have weak conjugate bases. • Weak acids have strong conjugate acids. (i) HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) acid base conjugate acid conjugate base (ii) CH3COOH(l) + H2O(l) H3O+ (aq) + CH3OOO-(aq) acid base conjugate acid conjugate base

  17. If HA is a strong acid in water, • HA is a successful donor of H+ in water • the reverse reaction hardly happens • A- is a poor acceptor of H+ • Ka (dissociation constant) is big HA + H2O H3O + + A- Equilibrium lies to the right. Strong acid , weak conjugate base Weak acid , strong conjugate base. Equilibrium lies to the left

  18. Common acids & conjugate bases in order of strengths

  19. Lewis theory • A Lewis acid is defined as a substance that can accept a pair of _________ from another atom to form a ________ covalent bond. • A Lewis base is defined as a substance that can __________ a pair of electrons to another atom to form a dative (coordinate) covalent bond. B: H+ +BH Lewis _____ Lewis ______

  20. Examples • Reaction between ammonia, NH3 and proton H3N: H+ +NH4 • Reaction between NH3 and BF3. H F H F H N B  H N B H F F H F F BF3 is a good Lewis ______ as there are _______electrons around the central boron atom which leaves room for 2 more electrons. Other common Lewis acid includes AlCl 3 and transition metal ions in aqueous solution.

  21. Reaction between a water molecule and proton H2O: H+ H3O+

  22. Lewis bonding In complex ions formed by transition metals The 6 water molecules, each donate a lone pair electrons from oxygen of their water molecules to the empty 3d orbitals of iron. What does each water molecule and iron(III) ion act as in the reaction above?

  23. Dative (Coordinate) bond • A dative covalent bond is always formed in a Lewis acid-base reaction. • For a substance to act as a base, it must have space to accept the lone pair of electrons.

  24. Strong and weak acids and bases Strong acid • When strong acid (HA) dissolves, virtually all acid molecules react with the water to produce hydronium ions (H3O+). HA + H2O(l)  H3O+(aq) + A-(aq) or HA  H+(aq) + A-(aq) 0% 100% 0% 100% Examples : HCl, H2SO4,HNO3, HClO4

  25. Strong and weak acids and bases Weak acid • When a weak acid dissolves in water, only a small % of its molecules (typically 1%) react with water molecules to release hydrogen or hydronium ions. The equilibrium lies on the left-hand side of the equation. HA + H2O(l) H3O+(aq) + A-(aq) or HA H+(aq) + A-(aq) 99% 1% 99% 1% Examples : CH3COOH, aqueous carbon dioxide

  26. Distinguish between strong and weak acids Base on the information above, how do we distinguish between strong and weak acids of the same concentration (e.g. HCl and CH3COOH)?

  27. How to distinguish between strong and weak acids? • A weak acid has a lower concentration of __________ and hence a higher _____ than a stronger acid of the same concentration. • A weak acid, because of its lower concentration of hydrogen ions, will have much poorer _____________ than a stronger acid of the same concentration. • Weak acids react more _______ with reactive metals, metal oxides, metal carbonates and metal hydrogencarbonates than strong acids of the same concentration. • Strong and weak acids can also be distnguished by measuring and comparing their enthalpies of neutralisation. What is the difference between the strength (strong and weak) and the concentrated (concentrated or dilute)?

  28. Strong and weak acids and bases Strong base • A strong base undergoes almost 100% dissociation/ionisation when in dilute aqueous solution. BOH B+(aq) + OH-(aq) 0% 100% Examples : NaOH, KOH, Ba(OH)2

  29. Strong and weak acids and bases Weak base • All bases are weak except the hydroxides of groups 1 and 2. • Weak bases are composed of molecules that react with water molecules to release hydoxide ions. In general for a weak molecular base, BOH • The equiibrium lies on the left side of the equation. BOH + (aq) B+(aq) + OH-(aq) Examples : aqueous ammonia, ethylamine, caffeine, bases of nuclei acids

  30. The pH indicator • scale that measures the strength of an acid and alkali. • pH of a substance is measured when it is dissolved in water. • pH stands for “power of hydrogen” • [H+] = 1 x 10-n moldm-3 ( n = pH number)

  31. The pH Scale

  32. pH probe and meter An accurate method of measuring pH value. A pH probe is dipped into the solution being tested and the pH value is then read directly from the meter.

  33. pH Calculation • pH is a measure of the concentration of H+ ions in a solution. • pH = -log10[H+(aq)] Example: If the concentration of H+ is 2.50 x 10-3 moldm-3 , what is the pH? pH = -log (2.50 x 10-3) = 2.60

  34. Example: Calculate the concentration of H+ of a solution that has a pH = 3.2. -log[H+] = 3.2 log[H+] = -3.2 [H+] = 6.31 x 10-4

  35. Example: Calculate the concentration of H+ and hence the pH of a 1.00 x 10-3 moldm-3 NaOH [OH-] = 1.00 x 10-3 moldm-3 [H+] x OH-] = 1.00 x 10-14 [H+] =

  36. Example: • (a) What is the pH of 10cm3 of 0.1 moldm-3 HCl? pH = -log (0.1) = 1 (b) If 90cm3 of water is added to the acid, what happens to the pH? Total volume = 100cm3 In 10cm3 solution, concentration of H+ is 0.1 moldm-3 • In 100cm3 solution, concentration of H+ is 0.01 moldm-3 • pH = -log (0.01) = 2 • (c) If the solution from (b) is diluted by a factor of 105 , what is the approximate pH? • The pH will increase to 6.

  37. Buffer • A buffer resists changes in _____ when small amounts of acid and alkali are added to it.

  38. Acidic Buffer • An acidic buffer solution can be made by mixing a weak ______ together with the _______ of the weak acid and a strong _______. (1) CH3COOH(aq) H +(aq) + CH3COO-(aq) (2) CH3COONa(aq) Na+(aq) + CH3COO-(aq)

  39. Acidic Buffer • If an acid is added, the extra H+ from the acid react with the excess ethanoate ions in (2) and are removed from the solution as ethanoic acid molecules (these have no effect on the pH). Hence the pH stays the same. CH3COO-(aq) + H +(aq) CH3COOH(aq) new

  40. Acidic Buffer • If an alkali is added, the OH- from the alkali react with the hydrogen ions from (1) removing them from the right hand side. There is, however, a large reservoir of ethanoic acid on the left hand side of this equilibrium able to dissociate and make more hydrogen ions, restoring the pH. CH3COOH(aq)+OH-(aq) CH3COO-(aq)+H2O(l)

  41. Acidic Buffer Acidic buffers are often made by taking a solution of a strong base and adding excess weak acid to it, so that the solution contains the salt and the unreacted weak acid. CH3COOH(aq)+NaOH(aq) CH3COONa(aq)+H2O(l) +CH3COOH(aq) limiting agent salt excess weak acid buffer

  42. Alkali Buffer • An alkali buffer with a fixed pH greater than 7 can be made from a weak base together with the salt of the base with a strong acid. • E.g. Ammonia and ammonium chloride NH4Cl(aq)  NH4+(aq) +Cl-(aq) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

  43. Alkali Buffer (1) NH4Cl(aq)  NH4+(aq) +Cl-(aq) (2) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) If H+ ions are added they will combine with OH- ions from (2) to form water and more of the ammonia will dissociate to replace them Adding more OH- ions that can react with the free ammonium ions (from equilibrium 1) producing more ammonia (as in equilibrium 2) and effectively being removed from the system. The ammonia molecules have no effect on pH an therefore the pH remains the same. In both cases, the hydroxide ion concentration and the hydrogen ion concentration remain constant.

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