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 - Dissociation constant (K a ) - Enthalpy of neutralisation

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

7. What is a base/alkali? A base is a substance like ___________ and ______________that react 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 _______. metal oxide metal hydroxide hydroxide soluble

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 _________ compound. When dissolves in water, it forms HCl acid which dissociate to form ions. hydrogen ions dissociate covalent

9. What is basicity (proticity)? Basicity refers to the no.of _________ atoms in one molecule of acid that can be replaced by a ______. E.g. HCl (monobasic), H 2 SO 4 (dibasic), H 3 PO 4 (tribasic) hydrogen metal

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 ________. donor acceptor

11. Types of acids  Acids that have single proton to donate – ___________. E.g. HCl(aq), HNO 3 (aq), HNO 2 (aq)  Acids that have 2 protons to donate – __________. E.g. H 2 SO 4 (aq), H 2 SO 3 (aq), H 2 CO 3 (aq)  H 3 PO 4 (aq) is _________. monobasic dibasic tribasic

12. Hydrogen chloride gas dissolved in water (solvent) HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq) The equation can be split into (i) HCl(aq) Cl - (aq) + H + (aq) (ii) H 2 O(l) + H + (aq) H 3 O + (aq) Acidic behaviour is a transfer reaction in different solvents. acid conjugate base baseconjugate acid

13. CH 3 COOH(l) + H 2 O(l) H 3 O + (aq) + CH 3 OOO - (aq) donates H + acidbase Acid-base conjugate pair conjugate base conjugate acid NH 3 (g) + H 2 O(l) NH 4 + (aq) + OH - (aq) Water is sometimes described as _______________ because it can accept or donate a proton. donates H + amphiprotic

14. Acid Strength & pK a Acid strength is the tendency of an acid to donate a proton. The more readily a compound donates a proton, the stronger is an acid.

15. Acidity Acidity is measured by an equilibrium constant, K eq. When a Bronsted-Lowry acid H-A is dissolved in water, an acid- base reaction occurs, and an equilibrium constant can be written for the reaction. H H-A + H-O-H A - + H-O-H : Keq= [products] [reactants] [H 3 O + ][A - ] [HA][H 2 O] =

16. Acidity and pK a The concentration of the solvent H 2 O is essentially constant, More convenient when describing acid strength to use “pK a ” values than K a. =[H 2 O]Keq [H 3 O + ][A - ] [HA] = Dissociation constant, K a

17. Competition between acid/base and its conjugate (i) HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq) acid base conjugate acid conjugate base (ii) CH 3 COOH(l) + H 2 O(l) H 3 O + (aq) + CH 3 OOO - (aq) acid base conjugate acid conjugate base (i) Water is a much stronger base than chloride ion and has a stronger tendency to accept _______. The equilibrium shifts more to the _______. (ii) Ethanote ion (CH 3 OOO - ) is a much stronger base than water molecule. The equilbrium shifts to the _______. protons right left

18.  Strong acids have weak conjugate bases.  Weak acids have strong conjugate bases. (i) HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq) acid base conjugate acid conjugate base (ii) CH 3 COOH(l) + H 2 O(l) H 3 O + (aq) +CH 3 OOO - (aq) acid base conjugate acid conjugate base

19. 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 +  K a (dissociation constant) is big HA + H 2 O H 3 O + + A - Equilibrium lies to the right. Strong acid, weak conjugate base Weak acid, strong conjugate base. Equilibrium lies to the left

20. Common acids & conjugate bases in order of strengths

21. Lewis theory  A Lewis acid is defined as a substance that can accept a pair of _________ from another atom to form a _______(coordinate) covalent bond.  A Lewis base is defined as a substance that can __________ a pair of electrons to another atom to form a dative covalent bond. B: H +  + BH Lewis _____Lewis ______ electrons dative donate acidbase

22. Lewis theory  Boroader definition that include compounds that do not have protons but exhibit acid/base behaviour. E.g. AlCl 3 (aq) + Cl - (aq) AlCl 4 - (aq) Lewis acid Lewis base Electron rich species react with electron poor species. All Bronsted Lowry acids are also Lewis acids

23. Examples  Reaction between ammonia, NH 3 and proton H 3 N: H +  + NH 4  Reaction between NH 3 and BF 3. H F H F H N B  H N B H F F H F F BF 3 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. acid6

24.  Reaction between a water molecule and proton H 2 O: H +  H 3 O +

25. 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? Water acts as Lewis base. Fe(III) acts as Lewis acid

26. 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 _________of electrons. lone pair

27. 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 (H 3 O + ). HA + H 2 O(l)  H 3 O + (aq) + A - (aq) or HA  H + (aq) + A - (aq) 0% 100% 0% 100% Examples : HCl, H 2 SO 4,HNO 3, HClO 4

28. 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 ________ side of the equation. HA + H 2 O(l) H 3 O + (aq) + A - (aq) or HA H + (aq) + A - (aq) 99% 1% Examples : CH 3 COOH, aqueous carbon dioxide left

29. 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 CH 3 COOH)? 0.1 mol dm -3 HCl (aq)0.1 mol dm -3 CH 3 COOH (aq) [H + (aq)]0.1 mol dm -3 0.0013 mol dm -3 pH1.002.89 Electrical conductivityhighlow Relative rate of reaction with magnesium fastslow Relative rate of reaction with calcium carbonate fastslow

30. 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.  Due to the lower concentration of hydrogen ions, a weak acid has poorer ___________________ than a stronger acid of the same concentration (equimolar).  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)? H+H+ pH electrical conductivity slowly

31. Strong and weak acids and bases Example of a strong base BOH  B + (aq) + OH - (aq) 0% 100% Examples : NaOH, KOH, Ba(OH) 2 Strong acid/base  A strong acid/base undergoes almost 100% ____________________ into its ions when in dilute aqueous solution. [readily donates H + / OH - ] ionization / dissociation

32. Strong and weak acids and bases Weak base  All bases are weak except the hydroxides of groups _______ in the Periodic Table.  In general for a weak molecular base, BOH  The equiibrium lies on the _____ side of the equation. BOH + (aq) B + (aq) + OH - (aq) Examples : aqueous ammonia, ethylamine, caffeine, bases of nuclei acids 1 and 2 left

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

34. The pH Scale

35. 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.

36. pH Calculation  pH is a measure of the concentration of H + ions in a solution.  pH = -log 10 [H + (aq)] Example: If the concentration of H + is (a) 1.0 x 10 -3 moldm -3 (b) 1.0 x 10 -2 moldm -3 (c) 2.50 x 10 -3 moldm -3, what is the pH? Compare (a) & (b) 2.60

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

38. Example: Calculate the concentration of H + and hence the pH of a 1.00 x 10 -3 moldm -3 NaOH

39. Example: (a) What is the pH of 10cm 3 of 0.1 moldm -3 HCl? (b) If 90cm 3 of water is added to the acid, what happens to the pH? (c) If the solution from (b) is diluted by a factor of 10 5, what is the approximate pH? [(a)1,(b)2,(c)7]

40. Buffer  A buffer resists changes in _____ when small amounts of acid and alkali are added to it. pH

41. Acidic Buffer  An acidic buffer solution can be made by mixing a weak ______ together with the _______ of the weak acid and a strong _______. (1) CH 3 COOH(aq) H + (aq) + CH 3 COO - (aq) (2) CH 3 COONa(aq) Na + (aq) + CH 3 COO - (aq) acid saltbase

42. Acidic Buffer  If an acid is added, the extra H + from the acid react with the excess ethanoate ions in (2) and are _________ from the solution as ethanoic acid molecules (these have no effect on the pH). Hence the pH stays the same. CH 3 COO - (aq) + H + (aq) CH 3 COOH(aq) new removed

43. Acidic Buffer  If an alkali is added, the OH - from the alkali react with the _____________ 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. CH 3 COOH(aq)+OH - (aq)  CH 3 COO(aq)+H 2 O(l) hydrogen

44. 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 NH 4 Cl(aq)  NH 4 + (aq) + Cl - (aq) NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq)

45. Alkali Buffer (1) NH 4 Cl(aq)  NH 4 + (aq) + Cl - (aq) (2) NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq) If H + ions are added they will combine with _____ (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 ____(from 1) producing more ammonia (as in 2) and effectively being removed from the system. The ammonia molecules have no effect on ____ an therefore the pH remains the same. In both cases, the hydroxide ion concentration and the hydrogen ion concentration remain constant. OH - NH + pH