1. Chapter 17 Acids and Bases

2. Necessary Terminology to Begin Hydronium = H 3 O + (aq) or H + (aq) Hydroxide = OH - (aq) Neutral Water = Equal amounts of above

3. General Info Bases Taste Bitter to the taste React with our skin to form soap (so it feels soapy) React with oils and greases (so often used in cleaning products (ammonia) Can cause dyes to change color. Electrolytes Corrosive

4. General Info about Acids Taste sour Corrosive Electrolytes Attack skin by dissolving fatty acids.

5. Common Acids and Bases Can you name some?

6. Common Sulfuric (Batteries in Car) Nitric (Explosives and Fertilizers Production) Hydrochloric (Steel Industry (Pickling))  Also called Muriatic

7. Common Aluminum Hydroxide (Deodorant) Magnesium Hydroxide (Laxative) Ammonia (Cleaner) Cough Syrups (taste awful without flavoring)

8. What is an Acid and Base? 3 different definitions that describe what Acids and Bases are  Arrhenius Acids and Bases  Bronsted-Lowry Acids and bases  Lewis Acids and Bases

9. Arrhenius Background His theory defined an acid as any substance that when added to water increases the hydronium ion concentration Bases as any substance when added to water that increase the hydroxide ion concentration.

10. Bronsted Acid and Bases The Bronsted definition of an acid comes from a man from Denmark who made his proposal in 1923. His theory helped to overcome the shortcomings of the Arrhenius definition by allowing us to describe solutions which were not aqueous.

11. Bronsted Acids and Bases His definition has acids as hydrogen donors and bases as hydrogen acceptors.

12. In Summary (Acids) An Arrhenius acid generates hydronium ions in water. A Bronsted acid donates hydrogens

13. In Summary (Bases) An Arrhenius base generates hydroxide ions in water A Bronsted base accepts protons

14. Practice Identifying On the following slides, identify the acid and base (forward reaction) and then whether each acid/base definition works

15. Questions CH 3 COOH (aq) + H 2 O (L)  CH 3 COO - (aq) + H 3 O + (aq) HCl (aq) + NH 3 (aq)  NH 4 + (aq) + Cl- NH 3 (aq) + H 2 O (L)  NH 4 + (aq) + OH- (aq)

16. Amphoteric A substance that can act as an acid or a base Water and Ammonia are common examples

17. Nomenclature Throw as many hydrogen’s onto the anion/polyatomic ion and change the ending from ate to ic and ite to ous. With the exception the halogens (Add hydrochloric acid = HCl) Perchlorate (ClO 4 -1 ) becomes Perchloric Acid (HClO 4 )

18. Nomenclature Nitrate (NO 3 -1 ) becomes Nitric Acid (HNO 3 ) Sulfite (SO 3 -2 ) becomes Sulfurous Acid (H 2 SO 3 ) Also: When writing the formula, if it is an acid, the H is placed at the beginning to denote that the chemical is an acid.

19. Major Ideas Strong and Weak Acids and Bases Conjugate Acids and Bases Bond Strength Acid/Base Equilibrium

20. What are Strong Acids/Bases? Strong acids and bases fully ionize when placed in water. The Unionized from of the acid is not present. There is no equilibrium, Strong acids and bases are completion reactions in water.  HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq)

21. If you take more chemistry, you will need to know these There are 6 strong acids  Nitric Acid (HNO 3 )  Sulfuric Acid (H 2 SO 4 )  PerChloric Acid (HClO 4 )  HydroBromic Acid (HBr)  HydroChloric Acid (HCl)  HydroIodic Acid (HI)

22. Strong Acids Every one of those, when placed in water, will ionize and all you will have is Hydronium and the Anion.  HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq)  Acid + Water  Hydronium + Anion

23. Strong Bases A strong base is any Alkali Metal with a Hydroxide. Such as: Sodium Hydroxide (NaOH), Potassium Hydroxide (KOH),

24. Big Idea behind Strong Acids If 1,000,000 moleculesof HCl are placed into 1.0 L of water, the Concentration of Hydronium is equal to the concentration of the Strong Acid placed into the solution. The concentration of Hydronium is 1,000,000 molecules in 1.0 L in this example.

25. Conjugate Acids/Bases When an Acid gives up its proton, the molecule left (minus a hydrogen) is called a conjugate base. The conjugate base has a negative charge and, being negative, has the ability to attract a nearby hydrogen (which is positive) to bond to it.

26. Conjugate Acids/Bases When a Base receives a proton, the molecule (plus a hydrogen) is called a conjugate acid. The conjugate acid has a positive charge and is looking to give up its positive charge to another molecule.

27. Conjugate Acids/Bases Summarized Conjugate Acid  An acid that forms when a base gains a proton Conjugate Base  A base that forms when an acid loses a proton

28. Example Acid Base  HC 2 H 3 O 2 (aq) + H 2 O   C 2 H 3 O 2 - (aq) + H 3 O + (aq)  Conj. Base Conj. Acid

29. Example Base Acid NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) Conj. Acid Conj. Base

30. Practice Questions What is the conjugate Base of the following Acids? HClH 2 SO 4 HydroniumAmmonium (NH 4 + )

31. Practice Label the following as Acid/Base/CB/CA HF + H 2 O  F - + H 3 O + CH 3 O - +H 2 O  CH 3 OH +OH -

32. What are Weak Acids/Bases? Weak acids and bases partially ionize when placed into water. An equilibrium is established where the Hydrogens are fought over (who gets to have the hydrogen?). Acetic acid (HC 2 H 3 O 2 ) when placed into water partially ionizes.

33. Weak Acid/Base Example Weak Acid  HC 2 H 3 O 2 (aq) + H 2 O   C 2 H 3 O 2 - (aq) + H 3 O + (aq) Weak Base  NH 3 (aq) + H 2 O (l)   NH 4 + (aq) + OH - (aq)

34. How weak is weak? If 100 weak acid molecules were put into a solution of water, only about 5 would react. Most weak acids are found with their hydrogen

35. New Ideas The pH Scale Calculating the Hydronium Concentration

36. pH Scale: Before we start… Quick Math: The Logarithm Scale On a logarithmic scale, a change in 1 represents a change in 10, a change in 2 represents a change of 100.  The Richter Scale (Earthquakes) An earthquake that registers a 5.0 is 1000x stronger than an earthquake that registers 2.0

37. The pH Scale Is a measure of how many Hydroniums are in the water. The pH means: Powers of Hydrogen  On the Log Scale A pH of 7 means the concentration of Hydronium is 0.0000001 or 1x10 -7

38. pH Scale pHpOH[H 3 O + ][OH - ] 01411 x10 -14 1130.1 1x10 -13 2120.011x10 -12 3111x10 -3 1x10 -11 4101 x10 -4 1 x10 -10 591 x10 -5 1 x10 -9

39. pH Scale pHpOH[H 3 O + ][OH - ] 681 x10 -6 1 x10 -8 771 x10 -7 1 x10 -7 861 x10 -8 1 x10 -6 951 x10 -9 1 x10 -5 1041 x10 -10 1 x10 -4 1131 x10 -11 1 x10 -3

41. Hydronium in Distilled Water pH: Why it adds up to 14 Water self ionizes (to a very small extent) 2 H 2 O (l) H 3 O + (aq) + OH - (aq) Kw = [H 3 O + ][ OH - ] Kw = 1.0 x10 -14 (This is a constant)

42. The pH scale When acids are added to water, they add to the hydronium concentration (decreasing the hydroxide concentration) When bases are added to water, they add to the hydroxide concentration (decreasing the hydronium concentration).

43. The pH Scale at neutral Kw = [OH][H] Kw = (1E -7 )(1E -7 ) = 1E -14 pH = 7 A pH of 7 is considered Neutral (it contains just as much base as acid)

46. Question Can you have a pH less than 0 and greater than 14? Yes, a 10.0 Molar HCl solution has a pH of -1

47. What is the pH? A solution has a hydronium concentration of 0.001? A solution has a hydroxide concentration of 0.00001? A solution has equal concentrations of hydronium and hydroxide?

51. Carbonic Acid Our Blood Buffer Carbon Dioxide in blood  Reacts with water forming Carbonic Acid  Same as pop More Carbon Dioxide = More Acid = Lower pH After a marathon, many runners have lower pH’s then 7.4. This is because of high carbon dioxide concentrations in the blood.

53. Breathing Rate When you hyperventilate, what happens to the blood pH? When you breath into a brown bag, what happens to the blood pH?

54. Neutralization Acids and Bases react to neutralize each other  Produce a salt and water HCl + NaOH  Water and NaCl

55. What is a Titration? A way to determine the acidity or baseness of a solution  You have an unknown solution, you can use a known concentration of acid to find out what the base concentration is

56. What is a Titration? Let’s say a student puts an unknown amount of NaOH into some water. You can determine how much NaOH was added by titrating the solution against a known concentration of acid

57. Equivalence Point At the equivalence point, equal amounts of acid and base have been added together, so there are no reactants left. In the case of adding HCl (a strong acid) to NaOH (a strong base), at the equivalence point, there is only water and salt. HCl + NaOH  NaCl + H 2 O

59. Indicators Indicators change color at different pH’s and are useful for knowing what the pH of the solution is (they won’t give exact, but they will let you know more or less acidic than some number)

60. During a titration, you place indicators into the solution to let you know when the pH has crossed a certain point. Different titrations have different equivalent points, so choosing an appropriate one is important.

61. For a SA and SB reaction, indicator isn’t important. For WA and WB reactions, it is more important

63. Phenolphthalein The two main differences are the extra H (in acidic solution) in the top left of the molecule, and the bonding of the central carbon