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IB Topic 8: Acids & Bases 8.1 Theories of Acids & Bases

IB Topic 8: Acids & Bases 8.1 Theories of Acids & Bases. 8.1.1 Define acids and bases according to the Brǿnsted-Lowry and Lewis theories. 8.1.2 Deduce whether or not a species could act as a Brǿnsted-Lowry and/or a Lewis acid or base.

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IB Topic 8: Acids & Bases 8.1 Theories of Acids & Bases

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  1. IB Topic 8: Acids & Bases8.1 Theories of Acids & Bases 8.1.1 Define acids and bases according to the Brǿnsted-Lowry and Lewis theories. 8.1.2 Deduce whether or not a species could act as a Brǿnsted-Lowry and/or a Lewis acid or base. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) 1

  2. Terms to Know • Brønsted-Lowry acid & base • Lewis acid & base • Strong acid & base • Weak acid & base • pH scale • Buffer • Conjugate acid & base • Indicator

  3. 8.1.1 Define acids and bases according to the Brǿnsted-Lowry and Lewis theories. Historical Aspects Acid: Comes from Latin word acidus, meaning sour or tart. Base: Comes from an old English meaning of the word, “to bring low.” When bases are added to acids, they lower the amount of acids. In the 1880’s Svante Arrhenius defined acids & bases: Acids are substances that, when dissolved in water, increase the [H+]. Bases are substances that, when dissolved in water, increase the [OH-]. 3

  4. Arrhenius definition Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water 4.3

  5. 8.1.1 Define acids and bases according to the Brǿnsted-Lowry and Lewis theories. Brǿnsted-Lowry Acids & Bases In 1923, Johannes Brǿnsted & Thomas Lowry defined acids & bases: Acids are substances that can transfer a proton to another substance (proton donor) Bases are substances that can accept a proton (proton acceptor) Lewis Acids & Bases Gilbert Lewis defined acids and bases: Lewis acid is an electron-pair acceptor Lewis base is an electron-pair donor 5

  6. A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor base acid

  7. Lewis Definition Lewis acid - a substance that accepts an electron pair Lewis base - a substance that donates an electron pair

  8. 8.1.2 Deduce whether or not a species could act as a Brǿnsted-Lowry and/or a Lewis acid or base. HCl is an acid HCl(g)  H+(aq) + Cl-(aq) (Arrhenius) Forms H+ ions in aqueous solution HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq) (Br-Lowry) H3O+(aq) is called the hydronium ion HCl donates a proton to water H+ accepts an electron pair from H2O (Lewis acid) H O H + H 8

  9. NaOH & NH3 are bases NaOH(s)  Na+(aq) + OH-(aq) (Arrhenius) Forms OH- ions in aqueous solution NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) (Br-Lowry) NH3 accepts a proton from water NH3 donates an electron pair to H+ (Lewis base) H H N H + H 8.1.2 Deduce whether or not a species could act as a Brǿnsted-Lowry and/or a Lewis acid or base. 9 9

  10. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) In any acid-base equilibrium, the forward reaction and the reverse reaction involve proton transfer: HX(g) + H2O(l)  H3O+(aq) + X-(aq) In the forward reaction, HX acts as an acid (donates a proton) and H2O acts as a base (accepts a proton). In the reverse reaction, H3O+ acts as an acid (donates a proton) and X- acts as a base (accepts a proton). HX & X- are an acid-conjugate base pair. H2O & H3O+ are a base-conjugate acid pair. They differ only in the presence or absence of a proton. 10

  11. Conjugate Pairs

  12. A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor conjugatebase conjugateacid base acid

  13. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) HNO2(aq) + H2O(l)  H3O+(aq) + NO2-(aq) acid base conj acid conj base A B-L acid loses a proton (H+) to form the conjugate base. A B-L base gains a proton (H+) to form the conjugate acid. remove H+ add H+ 13` 13

  14. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) H2O(l) + NH3(aq)  NH4+(aq) + OH-(aq) acid base conj acid conj base A B-L acid loses a proton (H+) to form the conjugate base. A B-L base gains a proton (H+) to form the conjugate acid. remove H+ add H+ 14` 14

  15. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) Consider the reaction: H2O(l) + HSO3-(aq)  H3O+(aq) + SO32-(aq) Identify the acid, the base, the conjugate acid and the conjugate base. 15` 15

  16. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) Consider the reaction: H2O(l) + HSO3-(aq)  H3O+(aq) + SO32-(aq) Identify the acid: HSO3-(aq) the base: H2O(l) the conjugate acid: H3O+(aq) the conjugate base: SO32-(aq) 16` 16

  17. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) What is the conjugate base for each of the following acids: HClO4; H2S; PH4+; HCO3-? 17` 17

  18. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) • What is the conjugate base for each of the following acids: • HClO4; H2S; PH4+; HCO3-? • The conjugate base for HClO4 is ClO4- since a B-L acid loses a proton (H+). • What is the conjugate acid of each of the following bases: CN-; SO42-; H2O; HCO3-? • The conjugate acid for CN- is HCN since a B-L base gains a proton (H+). 18` 18

  19. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) What is the conjugate acid for each of the following bases: HClO4; H2S; PH4+; HCO3-? 19` 19

  20. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) • What is the conjugate base for each of the following acids: • HClO4; H2S; PH4+; HCO3-? • The conjugate base for HClO4 is ClO4- since a B-L acid loses a proton (H+). • What is the conjugate acid of each of the following bases: CN-; SO42-; H2O; HCO3-? • The conjugate acid for CN- is HCN since a B-L base gains a proton (H+). 20` 20

  21. IB Topic 8: Acids and bases8.2: Properties of acids and bases 8.2.1 Outline the characteristic properties of acids and bases in aqueous solutions 21

  22. 8.2.1 Outline the characteristic properties of acids and bases in aqueous solutions Acids Sour taste Electrolytes Neutralization reactions with bases (such as metal hydroxides) to form a salt and water HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) Since HCl and NaOH are strong, they exist as ions so the reaction is: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  Na+(aq) + Cl-(aq) + H2O(l) Getting rid of “spectator” ions (ions that do not take part in the reaction), the equation can be written: H+(aq) + OH-(aq)  H2O(l) and the ΔH = -57.3 22

  23. 8.2.1 Outline the characteristic properties of acids and bases in aqueous solutions Acids Neutralization reactions with metal oxides to form a salt and water. H2SO4(aq) + CaO(aq)  Reactions with carbonates and hydrogen carbonates (bicarbonate) producing a salt, carbon dioxide and water. H2SO4(aq) + CaCO3(aq)  ? HCl(aq) + NaHCO3(aq)  ? 23

  24. 8.2.1 Outline the characteristic properties of acids and bases in aqueous solutions Acids Reaction with metals above hydrogen in the reactivity series to form a salt and hydrogen gas. H2SO4(aq) + Ca(s)  ? or H+(aq) + Ca(s)  ? Reactions with indicators. Indicators are substances that change color when the concentration of hydrogen ions changes. Ex. phenolpthalein, bromothymol blue, litmus H+ + In- HIn In- is one color, HIn is a different color. LeChatelier’s Principle Universal indicator is a mixture of different indicators and produces a range of colors. 24

  25. 8.2.1 Outline the characteristic properties of acids and bases in aqueous solutions Bases Bitter taste Feel slippery due to reaction with the oils on your skin. This forms a soap. Neutralization of acids (see acid properties). Displacement of ammonia (NH3) from ammonium salts. NH4Cl(s) + NaOH(aq)  ? or NH4+(aq) + OH-(aq)  ? Reactions with indicators. Bases do not react with indicators but the addition of a base changes the [H+] which affects the indicator. H+(aq) + OH-(aq)  H2O(l) 25

  26. IB Topic 8: Acids and bases8.3: Strong and weak acids and bases 8.3.1 Distinguish between strong and weak acids and bases, in terms of the extent of dissociation, reaction with water and electrical conductivity. 8.3.2 State whether a given acid or base is strong or weak 8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases, using experimental data. 26

  27. 8.3.1 Distinguish between strong and weak acids and bases, in terms of the extent of dissociation, reaction with water and electrical conductivity. Strong acids and bases dissociate completely in water Acid: HX(aq) + H2O(aq) H3O+(aq) + X-(aq) Base: YOH(aq) Y+(aq) + OH-(aq) When strong acids and bases dissolve in water, the solution consists of almost entirely of ions with a negligible amount of molecules. Write the dissociation reactions for H2SO4 & Ba(OH)2 27

  28. Strong and Weak Acids/Bases STRONG ACID: HNO3 (aq) + H2O (l) ---> H3O+ (aq) + NO3- (aq) HCl, HNO3, H2SO4 and HClO4 are examples of strong acids that dissociate completely in water.

  29. CaO Strong and Weak Acids/Bases • Strong Base: 100% dissociated in water. • NaOH (aq) ---> Na+ (aq) + OH- (aq) • Other common strong bases include: • Group 1 hydroxides (ex. KOH and LiOH) and • Ba(OH)2

  30. 8.3.1 Distinguish between strong and weak acids and bases, in terms of the extent of dissociation, reaction with water and electrical conductivity. Weak acids and bases only partly dissociate in water. Acid: HX(aq) + H2O(aq) H3O+(aq) + X-(aq) Base: NH3(aq) + H2O(aq) NH4+(aq) + OH-(aq) When weak acids and bases dissolve in water, the solution consists of almost entirely of molecules with a negligible amount of ions. Kc (Ka or Kb) << 1 Write the dissociation reactions for H2CO3 & C2H5NH2 30

  31. Strong and Weak Acids/Bases • Weak acids are much less than 100% ionized in water. • Examples include: Carboxylic acids, such as acetic acid aka ethanoic acid (CH3COOH) and carbonic acid (H2CO3)

  32. Strong and Weak Acids/Bases • Weak base: less than 100% ionized in water • One of the best known weak bases is ammonia • NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq) • As well as amines (nitrogen-containing compounds)

  33. 8.3.1 Distinguish between strong and weak acids and bases, in terms of the extent of dissociation, reaction with water and electrical conductivity. Strong acids and bases and electrical conductivity Since strong acids and bases dissociate completely into ions, these solutions are excellent conductors of electricity. However, there can be weak solutions of strong acids and bases which lowers the ion concentration and lowers the conductivity. Weak acids and bases and electrical conductivity Since weak acids and bases dissociate very little, few ions are present in solution so these are weak conductors of electricity. 33

  34. 8.3.2 State whether a given acid or base is strong or weak Strong acids and bases completely dissociate into their ions in aqueous solutions. Weak acids and bases only slightly dissociate into their ions in aqueous solution 34

  35. 8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases, using experimental data. Discussion of the lab: Properties of Acids & Bases 35

  36. IB Topic 8: Acids and bases8.4: The pH scale 8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. 8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values. 8.4.3 State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+(aq)]. 8.4.4 Deduce changes in [H+(aq)] when the pH of a solution changes by more than one pH unit. 36

  37. The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion.Under 7 = acid 7 = neutral Over 7 = base

  38. pH of Common Substances

  39. 8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. pH is a measure of the [H+(aq)] in a solution. pH = -log [H+(aq)] In a neutral solution the [H+(aq)] = 1 x 10-7 so the pH = -log(1 x 10-7) = 7 Addition of an acid to water increases the [H+(aq)] so the pH < 7. Low pH is acidic. Addition of a base to water decreases the [H+(aq)] so the pH > 7 . High pH is basic. 39 39

  40. Calculating the pH pH = - log [H+] (Remember that the [ ] means Molarity aka mol dm-3) Example: If [H+] = 1 X 10-10pH = - log 1 X 10-10 pH = - (- 10) pH = 10 Example: If [H+] = 1.8 X 10-5pH = - log 1.8 X 10-5 pH = - (- 4.74) pH = 4.74

  41. 8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. pH is a measure of the [H+(aq)] in a solution. pH = -log [H+(aq)] pH < 7 is acidic pH = 7 is neutral pH > 7 is basic Is a solution with a [H+(aq)] = 1 x 10-5 acidic, basic or neutral? Find the pH: pH = -log 1 x 10-5 = 5 so it is acidic Is a solution with a [H+(aq)] = 1 x 10-10 acidic, basic or neutral? Find the pH: pH = -log 1 x 10-10 = 10 so it is basic 41 41

  42. Try These! pH = - log [H+] pH = - log 0.15 pH = - (- 0.82) pH = 0.82 pH = - log 3 X 10-7 pH = - (- 6.52) pH = 6.52 • Find the pH of these: • A 0.15 M solution of Hydrochloric acid • 2) A 3.00 X 10-7 M solution of Nitric acid

  43. 8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. You can also calculate the [H+(aq)] if you know the pH [H+(aq)] = inverse log(-pH) In a solution with a pH = 6, what is the [H+(aq)]? [H+(aq)] = inverse log(-6) = 1 x 10-6 Is a solution with a pH = 11, what is the [H+(aq)]? [H+(aq)] = inverse log(-11) = 1 x 10-11 43 43

  44. 8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. Comparing pH of Strong and Weak Acids. Strong acids dissociate completely. A 0.010M HCl solution has a [H+(aq)] = 0.010M so the pH = 2 Weak acids do not dissociate completely so their [H+(aq)] concentration will be less than that of an equal concentration of a strong acid. A 0.010M ethanoic acid (CH3COOH) solution has a [H+(aq)] = 0.00042M so the pH = 3.4 A 0.01M carbonic acid (H2CO3) solution has a [H+(aq)] = 6.6 x 10-5 so the pH = 4.2 44 44

  45. 8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values. 45 45

  46. 8.4.3 State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+(aq)]. • An increase in one pH unit decreases the [H+(aq)] by a factor of 10. • A decrease in one pH unit increases the [H+(aq)] by a factor of 10. 46 46

  47. 8.4.4 Deduce changes in[H+(aq)] when the pH of a solution changes by more than one pH unit. • An increase in three pH units decreases the [H+(aq)] by a factor of 1,000. • A decrease in five pH units increases the [H+(aq)] by a factor of 105. 47 47

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