1. Unit 7: Acids and Bases

2. Acids and Bases: The Basics Acid comes from the Latin word, acidus, which means “sour.”  Ascorbic acid: C 6 H 8 O 8 – Citrus fruits  Acetic acid: CH 3 COOH – Vinegar  Hydrochloric acid: HCl – Toilet bowl cleaners  Carbonic acid: H 2 CO 3 - Soda  Sulfuric Acid: H 2 SO 4 – Fertilizers

3. Strength of Acids Strong acid: an acid that dissociates completely in aqueous solution. They are considered strong electrolytes.  Examples: HCl, HNO 3, H 2 SO 4, HBr  Acids tend to produce H 3 O + (Hydronium ion = Hydrogen ion)  H 3 O + is the same thing as H + HCl + H 2 O  H 3 O + + Cl - HNO 3 + H 2 O  H 3 O + + NO 3 - H 2 SO 4 + H 2 O  H 3 O + + HSO 4 - HBr + H 2 O  H 3 O + + Br-

4. Strength of Acids Weak acid: an acid that does not dissociate completely in aqueous solution.  Examples: H 3 PO 4, CH 3 COOH, H 2 CO 3 H 3 PO 4 + H 2 O ↔ H 3 O + + H 2 PO 4 - CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 COO - H 2 CO 3 + H 2 O ↔ H 3 O + + HCO 3 -

5. Strength of Acids

6. Strength of Bases Bases have a bitter taste and a slippery feel.  Also known as “alkaline”  Sodium Hydroxide (NaOH) – drain cleaners  Sodium bicarbonate (NaHCO 3 ) – baking soda  Potassium carbonate (K 2 CO 3 ) – ashes

7. Strength of Bases Strong base: A base that completed dissociates in water and yields aqueous OH - ions. Weak bases: A base that does not produce a large number of hydroxide ions (does not dissociate completely). Strong BasesWeak Bases NaOH  Na + + OH - NH 3 + H 2 O ↔ NH 4 + + OH - Ca(OH) 2  Ca 2+ + 2OH - KOH  K + + OH -

8. Acids and Bases Strong acids and strong bases are considered strong electrolytes, because a large concentration of ions are produced.

9. Bronsted-Lowry: Acid and Bases In the Bronsted-Lowry definition, an acid is any chemical that donates a hydrogen ion, H + (H 3 O + ), and a base is any chemical that accepts a hydrogen ion. B.A.A.D - “ B ases A ccept, A cids D onate”

10. Consider what happens when hydrochloric acid is mixed with water: HCl donates a H + … …resulting in a 3rd hydrogen bonded to oxygen Hydrogen ion H 3 O + HCl behaves as an acid (proton donor) and water behaves as a base (proton acceptor). Acids, when dissolved in water, release hydrogen ions.

12. Ammonia behaves as a base by accepting a H + from water, which, in this case, behaves as an acid. Hydroxide ion Bases tend to increase the concentration of hydroxide ions.

13. Hydroxide ion

14. Amphoteric Amphoteric: A substance that is capable of acting as either an acid or a base  It acts as a base when combined with something more strongly acidic than itself.  It acts as an acid when combined with something more strongly basic than itself. Water has the ability to react with itself.

15. Keep in mind … Acid-base interactions are almost seen as a behavior. For instance, water can behave as a base and as an acid. An ammonium ion may donate a H + back to OH - to reform ammonia and water.

16. Forward and reverse acid-base reactions proceed simultaneously and can therefore be represented by the double arrow.

17. Identify the acid or base behavior for each participant in the reaction: H 2 PO 4 + H 3 O + ↔ H 3 PO 4 + H 2 O Forward:  H 2 PO 4 - accepts a H + to become H 3 PO 4 H 2 PO 4 - behaves as a base H 3 O + donates a H + to become H 2 O H 3 O + behaves as an acid

18. Identify the acid or base behavior for each participant in the reaction: H 2 PO 4 - + H 3 O + ↔ H 3 PO 4 + H 2 O Backwards:  H 3 PO 4 donates a H + to become H 2 PO 4 - H 3 PO 4 behaves as a acid H 2 O accepts a H + to become H 3 O + H 2 O behaves as a base

19. Conjugate Acid-Base Pairs In any acid-base equilibrium (↔), it involves the transfer of H +. An acid and a base that only differs in the presence or absence of a H + are called a conjugate acid-base pair.  Every acid has a conjugate base For example: H 2 O (acids donate) can become OH -  Every base has a conjugate acid For example, H 2 O (bases accept) can become H 3 O +

20. Conjugate Acid-Base Pairs HNO 2 (aq) + H 2 O (l) ↔ NO 2 - (aq) + H 3 O + (aq) AcidBase Conjugate Acid Conjugate Base donates H+ Accepts H+ HNO 2 donates an H + and becomes its conjugate base, NO 2 - H 2 O accepts an H + and becomes its conjugate acid, H 3 O +

21. Conjugate Acid-Base Pairs HCl + H 2 O H 3 O + + Cl - acid Conjugate base Conjugate acid Accepts H + Donates H +

22. Acid Dissociation Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 280 HCl Conjugate base Acid Conjugate pair + 1- Cl H

23. Conjugate Acid-Base Pairs NH 3 + H 2 O NH 4 + + OH - base Conjugate acid acid Conjugate base Accepts H + Donates H +

24. What is the conjugate base of each of the following acids: Step 1: Remember that acids DONATES an H + Step 2: Its conjugate base always has an extra negative charge HClO 4  ClO 4 - H 2 S  HS - HCO 3 -  CO 3 2-

25. What is the conjugate acid of the following bases? Step 1: Remember that all bases ACCEPT an H + Step 2: Its conjugate acid always has one LESS negative charge CN -  HCN H 2 O H3O+H3O+ HCO 3 -  H 2 CO 3

26. Neutralization Reactions Neutralization reactions occur between an acid and a base.  These reactions often produce a salt, created from the positive ion of the base and the negative ion from the acid. ACID BASE SALT HCN + NaOH  NaCN + H 2 O HNO 3 + KOH  KNO 3 + H 2 O

28. Neutralization Reactions Predict the salt that is produced in the following: 2 HCl + Ca(OH) 2  ____________ + 2 H 2 O HF + NaOH  ____________ + H 2 O CaCl 2 NaF Prevents tooth decay De-ice roads

29. What is pH?

30. In this reaction, a water molecule gains a H + and the second water molecule must lose a H +. In pure water, the number of H + = the number of OH -  The concentration of H + and OH - each is extremely low – about 1 x 10 -7 M

31. What is pH? [H 3 O + ][OH - ] = K w [1.0 x 10 -7 ][1.0 x 10 -7 ] = K w 1.0 x 10 -14 = K w The dissociation constant of water, K w, means that no matter WHAT is dissolved in water, the product of H + and OH - always equals 1.0 x 10 -14

32. What is pH? [H 3 O + ][OH - ] = K w = 1.0 x 10 -14 Pure Water [1.0 x 10 -7 ][1.0 x 10 -7 ] = K w = 1.0 x 10 -14 HCl Added [1.0 x 10 -5 ][1.0 x 10 -9 ] = K w = 1.0 x 10 -14 If a small amount of HCl is added to water, it dissociates and increases the H + from 1.0 x 10 -7 to 1.0 x 10 -5. Therefore, the OH - concentration decreases so that the product of H + and OH - is still equal to K w.

33. Sample Problem What is the concentration of H + ions if the concentration of OH - ion is 1.0 x 10 -3 M? [H 3 O + ][OH - ] = K w [H 3 O + ][1.0 x 10 -3 M] = 1.0 x 10 -14 M [H 3 O + ] = [H 3 O + ] =1.0 x 10 -11 M 1.0 x 10 -14 M 1.0 x 10 -3 M

34. What is pH? In an acidic solution, [H 3 O + ] > [OH - ] In a basic solution, [H 3 O + ] < [OH - ] In a neutral solution, [H 3 O + ] = [OH - ]

36. Sample Problem How does adding ammonia, NH 3, to water make a basic solution when there are no hydroxide ions in the formula for ammonia? NH 3 + H 2 O  NH 4 + + OH - Ammonia increases the OH- concentration, thereby lowering the H + concentration. Because [H+] < [OH-], the solution is basic.

37. What is pH? The pH scale is a numeric scale used to describe acidity. pH = -(log[H 3 O + ]) Consider a neutral solution, [H + ] = 1.0 x 10-7 M pH = -(log [1.0 x 10 -7 ]) pH = -(-7) pH = 7

38. What is pH? Acidic solutions have greater H + concentrations, which lowers its pH. Acidic solutions: pH < 7 Consider [H + ] = 1.0 x 10 -4 M pH = -(log [H + ]) pH = -(log [1.0 x 10 -4 M]) pH = -(-4) pH = 4

40. What is pH? Basic solutions have pH values > 7, because its H + concentrations are less. Consider [H + ] = 1.0 x 10 -8 M pH = -(log [H + ]) pH = -(log [1.0 x 10 -8 M]) pH = -(-8) pH = 8

41. What is pH?

42. Rainwater is Acid & Ocean Water is Basic

43. Rainwater is Acidic: pH 5-6 The source of this acidity is carbon dioxide, the same gas that gives fizz to soda pop.  There are 760 billion tons of CO 2 in the atmosphere that undergo this reaction: CO 2 (g) + H 2 O (l)  H 2 CO 3 (aq) Carbonic acid Carbonic acid lowers the pH, accelerating the erosion of land and historical artifacts.

45. What is classified as ACID RAIN? Acid rain: rain that has a pH < 5.  Source: Airborne pollutants that are absorbed by atmospheric moisture, most commonly – sulfur dioxide.  Sulfur dioxide is readily convert to sulfur trioxide… 2 SO 2 (g) + O 2 (g)  2 SO 3 (g)  …which reacts with water to form sulfuric acid. SO 3 (g) + H 2 O (l)  H 2 SO 4 (aq) Acid rain affects vegetation and ecosystems.

47. ACID RAIN CYCLE

48. Impact of Acid Rain: Midwest Midwest: The ground contains calcium carbonate (basic), which often neutralizes the acid rain before much damage is done.

49. Liming In order to rains the pH of acidified lakes and rivers by adding calcium carbonate – a process called liming. Long-term solution: Prevent the sulfur dioxide and other pollutants from entering the atmosphere in the first place.  Shift fossil fuels to nuclear and solar energy

50. Impact of Acid Rain: Northeast Northeast: The ground contains LITTLE CaCO 3. The effect of acid rain on lakes and rivers accumulate.

51. Buffer Solutions Resist Change in pH

52. Buffer solution Buffer solution: any solution that resists change in pH  (1) neutralizes any added base  (2) neutralizes any added acid This does not mean that the pH remains unchanged, it just resists LARGE changes in pH

53. Example of Buffers Blood: optimal pH of 7.35 to 7.45  Primary buffer system combines (1) carbonic acid and (2) sodium bicarbonate  Carbon dioxide in the blood stream reacts with water to produce carbonic acid CO 2 + H 2 O  H 2 CO 3 We fine-tune the levels of carbonic acid in our blood by breathing!

54. Aspirin overdose: Alkalosis Aspirin (acetylsalicylic acid) is an acid chemical that when taken in large amounts can overwhelm the blood buffering system, dropping the blood pH. Symptom: Hyperventilate  Exhaling at excessive rates is your body’s attempt to lower the concentration of carbonic acid.  Alkalosis: suddenly raising your blood pH can be life- threatening  Acidosis: (if the breathing rate is too slow) suddenly lowering your blood pH (increased H 3 O + concentration)

55. Alkalosis If our breathing becomes too fast (hyperventilation)… Carbon dioxide is removed from the blood too quickly. This accelerates the rate of degradation of carbonic acid into carbon dioxide and water. The lower level of carbonic acid encourages the combination of hydrogen ions and bicarbonate ions to make more carbonic acid. The final result is a fall in blood H 1+ levels that raises blood pH which can result in over-excitability or death. Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291

56. Acidosis If breathing becomes too slow (hypoventilation)… …free up acid, pH of blood drops, with associated health risks such as depression of the central nervous system or death. The normal pH of blood is between 7.2 – 7.4. This pH is maintained by the bicarbonate ion and other buffers. Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291